Igcse2

IGCSE Chemistry-Dr. D. Bampilis

IGCSE CHEMISTRY

NOTES


1 Τhe particulate nature of matter

Atom: The smallest particle of matter

Molecule: A small particle made from more than one atom bonded together

Element: A substance made of only one type of atom

Compound: A substance made from two or more different elements bonded together

States of matter:

Solid:

1. Strong forces of attraction between particles

2. Have a fixed pattern (lattice)

3. Atoms vibrate but can’t change position therefore fixed volume and shape

4. Can’t be compressed

Liquid:

1. Weaker attractive forces than solids

2. No fixed pattern, liquids take up the shape of their container but have a fixed volume

3. Particles slide past each other.

4. Can’t be compressed

Gas:

1. Almost no intermolecular forces

2. Particles are far apart, and move quickly

3. They collide with each other and bounce in all directions.

4. Can be compressed

The Kinetic Theory of Matter States:

The kinetic theory is a theory put together by the finest chemists and physicians of all time. It

consists of a number of true facts related to matter and their states. The theory explains the

behavior of matter and their physical properties.

The kinetic theory of matter states:

All matter is made up of tiny, microscopic moving particles. And each matter has a different

type of particles with different size and mass.

Particles are in continuous movement. All particles are moving all the time in random

directions (Brownian motion).

The speed of movement depends on the mass of the particle, temperature and several other

factors that you will know later on.

Kinetic means movement, and so kinetic energy means movement energy.


Comparing Properties of Solids, Liquids and Gases:

Molecular

Structure

Solid Liquid Gas

Particles

Arrangement

Very closely packed

Regularly arranged in

lattice

Closely packed

Irregular arrangement

Very far apart

Very irregular

arrangement

Intermolecular

Spaces

Almost none

Negligible

Minimal

Tiny spaces Very large

Intermolecular

Forces Extremely strong

Not weak

Weaker than in solids Very weak

Movement of

Particles

Vibrating in a fixed

position

Slowly slide over each

other randomly

Fast movement in

random direction

Shape Fixed definite shape

No fixed shape

Depends on the

container

No fixed shape

Volume fixed fixed

No fixed volume

Depends on the

container

Compressibility Cannot be compressed Can be hardly

compressed Very compressible

Diffusion Cannot diffuse Diffuses slowly Diffuses quickly

• melting – freezing – boiling – condensing – subliming – desubliming


Diffusion is the random movement of liquid or gas particles to fill the available space and

spread evenly. For instance, if you pass by a trash can, you can smell the ugly scent of trash.

This is because molecules from the garbage diffused out of the can to the air which you

breathed in.

Diffusion rate depends on several factors, these are:

Mass of the substance. The lighter the substance (lower Mr or Ar) the faster it diffuses

Temperature. The more kinetic energy the particles have, the faster they move and diffuse.

Presence of other substance. Diffusion is faster when it occurs in an area where there are

fewer particles of other substances present. This is why diffusion is extremely fast in vacuums.

This is because the diffusing particles have less other particles to stand in their way.

Intermolecular spaces. This is why gases diffuse faster than liquids and solids do not diffuse.

TIPS ON SPECIFIC TOPICS

• Remember that most of the particles in liquids are touching one another. It is a common error

to think that they are well separated.

•Diffusion is due to the random movement of particles so they spread out everywhere. In an

exam, try not to give an answer involving movement of particles from high to low

concentration as this suggests that the particles know where they are going.


2 Experimental techniques

2.1 Αpparatus for the measurement Time(s): Stopwatch Clock

Temperature(0C): Thermometer

Mass(kgr-gr): Balance

Volume(m3-dm

3-cm

3)of liquids:

beaker -burette -pipettes -measuring cylinder -volumetric flask

of gases:

gas syringe-upturned measuring cylinder

2.2 (a) Criteria of purity

• Paper chromatography is a technique that can be used to separate mixtures of dyes or

pigments and is used to test the purity of a mixture or to see what it contains

•different solubility in the solvent

• different degrees of attraction for the filter paper

• The importance of purity in substances in everyday life, e.g. foodstuffs and drugs

Purity can be measured in a number of ways:

• melting point/boiling point (impurities increase b.p. and decrease m.p.)

• chromatography

• Rf values

.

Rf = distance moved by the compound ÷ distance moved by the solvent

• Chromatography techniques can be applied to colourless substances by exposing

chromatograms to substances called locating agents(ninhydrin for aminoacids)


2.2 (b) Methods of purification

• Evaporation used to separate a solid from a solution.

• Filtration to remove solid particles from a liquid.

• Crystallization removes the solvent, to leave the solute.

• Distillation used to separate a solvent from a solution.

1. Salty water is heated

2. The water vapour cools in the condenser and drips into a beaker

3. The water has condensed and is now in the beaker, the salt stays behind


• Fractional distillation used to separate liquids from each other, produces a number of

substances from the original mixture (e.g. petroleum).

1. Water and ethanol solution is heated

2. The ethanol evaporates first, cools, then condenses

3. The water left evaporates, cools, then condenses

• Sedimentation allows an insoluble solid to separate out and sink to the bottom of a container


• Centrifugation a spinning motion increases the force of gravity that quickly separates a solid

from a suspension.

• Decanting pouring off liquid (e.g. pouring off excess water from a pot of peas).

• Magnetic Separation a method for separating one solid (usually iron) from a mixture of solids,

very useful for separating aluminium cans from steel cans.

• Solvent extraction used to separate two solutes dissolved in a solvent.

• Chromatography used to separate different substances from a solution



•When measuring out volumes, think about the accuracy needed. A burette or volumetric

pipette is far more accurate than a measuring cylinder.

•When drawing chromatography apparatus, you must draw the origin line on the chromatogram

so that it is above the starting level of the solvent.

•Remember that pure substances have definite sharp melting points and boiling points. Impure

substances melt and boil over a range of temperatures.

•When describing crystallization, the answer ‘heat the solution’ is not enough. You need to

write ‘evaporate off some of the water and then leave to cool’.

•When choosing a method to purify a mixture, think about the states and solubilities of the

substances in the mixture.

• If you are distilling an aqueous solution of a salt, the salt itself does not evaporate as it has too

high a boiling point. Only the water evaporates


3 Atoms, Elements and Compounds

3.1 Atomic structure and the Periodic Table

• Protons, neutrons and electrons

• Electrons shells / energy levels

• Proton/atomic number and nucleon/mass number

• Periodic Table

• Isotopes

• Types of isotopes as being radioactive and non-radioactive

• Medical and industrial uses of radioactive isotopes

• Electron configuration

• Period(number of shells)


• Group(number of e in outer shell)

• Valence electrons – chemical properties- group

• Noble gases electronic structures

3.2 Bonding: the structure of matter

• Element a substance that cannot be split into anything simpler, in a chemical reaction. Each

element has a unique proton number.

• Mixture two or more elements mixed together BUT that are not chemically combined

• Compound a substance in which two or more different elements are chemically combined

(molecular – ionic).

Differences compound mixture

are pure substances impure substances

made up of two or more elements combined

chemically

substances mixed

physically

composition a fixed ratio varying ratios

properties fixed , different from its

constituents

no fixed ,same of its

constituents.

can be separated only by chemical

methods

easily by physical

methods

• Density

Differences in

Physical Properties

Metal Non-metal

conductors of heat

and electricity

good poor

malleable - ductile yes no - brittle

lustrous yes - shiny no -dull

at room temperature solids (exception is mercury) solids or gases(exception is

bromine)

melting point high(exception group I, Hg) low(exception C,Si)

density high(exception group I, Ga) low(most)

sonorous yes no

• Alloy, such as brass, a mixture of a metal with other elements


3.2 (a) Ions and ionic bonds

• Formation of ions by electron loss(cations) or gain(anions)

• Ionic bonds between metallic and non-metallic elements

Group

1

Group

2

Group

3

Group

4

Group

5

Group

6

Group

7

Group

0

Example

element Na Mg Al C N O Cl He

Charge + 2+ 3+ Note 1 3- 2- - Note 2

Symbol

of ion Na+ Mg

2+ Al

3+ Note 1 N

3- O

2- Cl

- Note 2

Note 1: Carbon and silicon in Group 4 usually form covalent bondsby sharing electrons.

Note 2: The elements in Group 0 do not react with other elements to form ions.

• Octet rule

• Lattice structure of ionic bond

The diagram shows part of the crystal lattice of sodium chloride:

Properties of Ionic Compounds:

Hard solids at room temperature,

High melting and boiling points because of strong attraction forces,

When solid they are electrical insulators but conduct electricity when molten or aqueous,

Water soluble.

3.2 (b) Molecules and covalent bonds

• Single covalent bonds in H2, Cl2 , H2O, CH4 and HCl as the sharing of pairs of electrons

leading to the noble gas configuration (non metal- non metal)

• Shared pair- lone pair

• Electron arrangement in more complex covalent molecules such as N2, C2H4, CH3OH and

CO2 , Double- triple bond

http://www.bbc.co.uk/education/guides/zd8hvcw/revision/1#glossary-z6bhb9q


How many bonds?

Element Number of bonds

Group 4 Carbon 8 - 4 = 4

Group 5 Nitrogen 8 - 5 = 3

Group 6 Oxygen 8 - 6 = 2

Group 7 Chlorine 8 - 7 = 1

Hydrogen forms one covalent bond.

The noble gases in Group 0 do not form any.

• Intermolecular forces: weak – m.p. , b.p.

• Valency of an atom: the number of electrons that would be gained, lost or share if it reacts

with other atoms.

Types of Covalent Structures:

There are two types of covalent structures:

Simple Molecular Structure

Giant Molecular Structure

http://chemistry.about.com/od/chemistryglossary/a/atomdefinition.htm


Simple Molecular Structure:

They are simple and contain only a few atoms in one molecule. Covalent bonds between the

atoms within a molecule (intermolecular bonds) are strong but they have weak bonds between

molecules (intermolecular bonds). These forces increase as the size of the molecule increases.

Giant Molecular Structure:

They are also known as macromolecular structures. One molecule contains hundreds of

thousands of atoms. They have extremely strong bonds between the atoms (intermolecular

bonds).

Properties of Covalent Compounds:

Simple molecular structures are usually gases or liquids and sometimes solids with low

melting points; this is because of weak forces of attraction between the molecules which can be

broken easily.

Giant molecular structures have very high melting points because the whole structure is held

together with very strong covalent bonds.

Most of them do not conduct electricity

Most of them are insoluble in water

Differences in Chemical

Properties

Metal Non-metal

electrons in the outer

shell

1-3(exception is

hydrogen)

4-8

valence electrons lose easily gain or share

form cations anions(exception is

hydrogen)

form oxides basic acidic

react with acid form hydrogen no

Differences ionic compounds

covalent compounds

volatility high low

melting points and

boiling points

high low

solubility in water usually soluble the majority do not

dissolve

electrical conductivity molten or dissolved in

water

ionic solids are good

insulators

don't conduct electricity

in an aqueous solution

form crystal lattice molecules

are hard tend to be soft and

relatively flexible

3.2 (c) Macromolecules

• Giant covalent structures of graphite and diamond(allotropes of C)

What are allotropes? When an element exists in several physical forms of the same state, it is

said to exhibit allotropy. Each form of this element is an allotrope. Lots of elements exhibit

allotropy. Carbon has two very popular allotropes, diamond and graphite. Diamond and

graphite are both made of carbon only. However, they look very different and have different

physical properties. They are both giant molecular structures.


• Graphite: 3 covalent bonds

hexagon – layers(weak bonding)

delocalized e

conductor(electods)

high m.p., b.p.

used as a lubricant and in pencil leads(can flake off easily)

• Diamond : 4 covalent bonds

insulator

high m.p., b.p.

very hard

used for cutting and drilling

• Macromolecular structure of silicon (IV) oxide (silicon dioxide-sand - quartz)

Similarity in properties between diamond and silicon (IV) oxide, related to their structures

The graphic shows the molecular structure of graphite and diamond(two allotropes of

carbon) and of silica (silicon dioxide).

3.2 (d) Metallic bonding

• Metallic bonding is a lattice of positive ions in a ‘sea of electrons’

• The electrical conductivity, malleability and high m.p. and b.p. of metals

The electrical conductivity of a metal decreases with increasing temperature(the vibration of

cations inhibit the move of electrons)


•In an exam you will always be given a Periodic Table .You can use your Periodic Table to find

out the number of protons in an atom. You can also use it to calculate the number of neutrons.

•You did not need to know the details about radioactivity or about α-, β- or γ-radiation. Don’t

try to remember lots of uses for radioisotopes –just remember one medical and one industrial

use.

•Make sure that you can draw the electronic structure of the first 20 elements in rings

containing electrons. If you are simply asked ‘what is the electronic structure of sodium?’

•You should learn the definitions of elements, compounds and mixtures. You may be asked to

write these definitions in an exam.

•If you are asked how to tell the difference between a metal and a non-metal it is best to select

conductivity, malleability or ductility as properties. These have fewer exceptions to the general

rules.

http://www.bbc.co.uk/education/guides/z94xsbk/revision/4#glossary-zq7y87h


•When drawing the electronic structure for an ion, make sure that the charge of the ion is shown

at the top- right-hand corner just outside the square brackets. Do NOT put the charge in the

nucleus.

• When drawing dot-and-cross diagrams remember to pair up the bonding electrons in the

overlap area between the atoms. Don’t put them outside the area where the atoms join.

• When drawing the electronic structure of compounds with double and triple bonds, make sure

that you draw the atoms large enough so that all the bonding electrons can fit into the overlap

area of the atoms.

•Remember that compounds of metals with non-metals are likely to be ionic. Compounds of

non-metals with other non-metals are covalent.

•When explaining why graphite conducts electricity, make sure that you state that electrons in

the layers can move along. Do not write ‘The electrons move.’- that suggest that the electrons in

the covalent bonds can move through the structure as well.

•It is a common error to suggest that conduction in metals is due to moving ions. Remember

that it is only the delocalized electrons which move. The positive ions remain fixed in position

within the giant lattice.

•When writing symbols containing two letters, make sure that the second letter is a small one.

Cl is correct for chlorine. CL is wrong.

•Take care when writing the second atom in a formula. Co2 is not acceptable for carbon dioxide

and neither is H2o for water. The symbol for oxygen is always a capital O.

•When asked to write the formula of an ionic compound from a diagram of its structure, make

sure that you write the formula as the simplest ratio. For example, CaBr2 not a Ca8Br16.

• It is a common mistake to count the bonds and not the electrons when asked about the number

of electrons shared between the atoms in a molecule. For example, the number of shared

electrons in methane is eight not four

• Take care when writing electronic structures including hydrogen. Always show the hydrogen

atom either as a circle or (if ionic) by its symbol. It is best practice to write the symbol of the

atom in the centre so it is clear to the examiner which atom is which

• When writing dot-and-cross diagrams for ionic structures, put the charge outside of the

brackets, at the top, not in the centre of the atom

• When asked about the number of covalent bonds in a compound, focus on the outer energy

level / shell electrons that are shared, not the total number of electrons. Remember that some

molecules have non-bonding pairs of electrons e.g. nitrogen

• When drawing dot-and-cross diagrams for molecules such as nitrogen which have only three

bonding pairs of electrons, don’t forget to draw in the lone pairs of electrons. Remember that

there must be eight electrons surrounding each atom

• Practice drawing diagrams of giant molecule structures, including silicon dioxide, diamond

and graphite, as these are nearly always drawn badly. You must show the continuation bonds


4 Stoichiometry

• Symbols of the elements

• Formulae of simple compounds (group- valency-formula)

• Naming compounds

AXBY : A B ide

• Molecular formula – Empirical formula - Structural formula

• Deduce the formula of a simple compound from a model or a diagrammatic representation

Molecular compounds from structural formula

Ionic compounds from simplest ratio

• Determine the formula of an ionic compound from the charges on the ions present

The Periodic Table and Charges:

Group

(Charge)

1

(+1)

2

(+2)

Transition

metals

3

(+3)

4

(±4)

5

(-

3)

6

(-

2)

7

(-

1)

Ions

present

Li+

Na+

K+

Be2+

Mg2+

Ca2+

Ba2+

Cu2+

/

Cu+

Fe2+

/

Fe3+

Zn2+

Ag+

Al3+

C

Si

Pb2+

N3-

P3-

O2-

S2-

F-

Cl-

Br-

I-

Compound Ions:

Oxidation

State Name Symbol

+1 Ammonium Ion NH4+

-1

Hydroxide Ion

Nitrate Ion

Nitrite Ion

Manganate(VII) Oxide Ion

Hydrogen Carbonate Ion

OH-

NO3-

NO2-

MnO4-

HCO3-

-2

Carbonate Ion

Sulfate Ion

Sulfite Ion

Dichromate (Vi) Ion

CO32-

SO42-

SO32-

Cr2O72-

-3 Phosphate Ion

Phostphite Ion

PO43-

PO33-


• Word equations

• Simple balanced chemical equations

• Diatomic elements :H2, O2, N2, F2, Cl2 , Br2, I2

• Construct equations with state symbols

• Ionic equations

Substances form ions:

-metals/non metals

-acids

-ammonium compounds

Spectators ions

• Relative atomic mass, Ar

Calculating relative atomic mass from isotopic abundance

• Relative molecular mass, Mr , as the sum of the relative atomic masses

• Relative formula mass or Mr for ionic compounds

• Calculations involving reacting masses in simple proportions

4.1 The mole concept

• mole

• Avogadro constant

• Molar mass: the relative formula mass in g

n= or m=nM

n:number of mole

m: mass in g

M:molar mass

• Molar gas volume, taken as 24 dm3 at room temperature and pressure (20

0C – 1 atm)

• Solution concentrations expressed in g/dm3 and mol/dm

3

C= n=CV

C: concentration in mol/dm3

V:volume of solution in dm3

n:mole of solute


• Stoichiometry

• Limiting reactants

• Calculate stoichiometric reacting masses and volumes of gases and solutions

• Titration

• Percentage by mass of an element in a compound : using Ar and relative formula mass

• Calculate empirical formulae and molecular formulae

• % yield= 100

• % purity= 100


•When balancing symbol equations you must not change any of the formulae. Always balance

by putting large numbers in front of the formulae. For example, balancing CaΟ by making it

into CaO2 is wrong.

•When writing ionic equations, first identify the reactants or products that are not ionic. These

will be solids, liquids or simple molecules like chlorine. It is only then that you can separate the

other compounds into ions.

•If a formula has brackets, first work out the atomic masses inside the brackets then multiply by

the number outside. Finally, add the atomic masses which were not bracketed.

•When doing calculations put the relative formula masses or moles below the appropriate

reactants or products in the symbol equation so that you can see the reactants or products are

relevant. Be sure to take the stoichiometry of the equation into account.

•The limiting reactant is the reactant that is NOT in excess. It has the smaller number of moles.

Be careful though – you must also take into account the ratio in which the reactants combine.

•When out gas volumes first find the number of moles and then multiply this by 24. The answer

is then in dm3. Remember that the molar gas volume is given at the bottom of your Periodic

Table.

•Always show your working in calculations if a question is worth more than one mark. If you

make an error at the start – for example use an incorrect molar mass – you can still gain marks.

•When calculating empirical formulae, make sure that between steps 1 and 2 you don’t round

up the figures. This often leads to errors.

•Mole calculations involving concentrations are easier if you change cm3 to dm

3 and then use

the formula concentration= number of moles/ volume of solution in dm3

• If asked for a word equation, do not write a symbol equation. A word equation tests

knowledge of chemical names. Although a correct symbol equation is often accepted this is not

guaranteed and if you make an error, you won’t get the mark

• A common error is to think that a nitrate ion has a 2- charge. The formula for the nitrate ion is

NO3-.This makes the formula for nitric acid HNO3


• The charge on a silver ion is 1+. A common mistake is to think that silver has a 2+ charge

• When working out formulae, don’t be confused by oxidation numbers. A common mistake is

to think that the formula for lead(IV) oxide is PbO4 or that lead(II) nitrate is Pb2(NO3). In a

formula you have to balance the positive and negative charges. Lead(IV) = 4+, lead(II) = 2+,

oxide = 2- and nitrate = 1-. So lead(IV) oxide is PbO2, and lead(II) nitrate is Pb(NO3)2

• If asked to name a salt formed in a particular reaction, don’t put down any other product or

you will lose a mark

• When calculating moles, if you are given an equation such as:

Mg + 2CH3CO2H → (CH3CO2)2Mg + H2

ignore the 2 in the equation when calculating the molar mass of ethanoic acid. The molar mass

of ethanoic acid is 60, not 120. However, remember when calculating reacting masses that the 2

needs to be taken into account


5 Electricity and chemistry

• Electrolysis is a process in which electricity is used to break compounds down into their

elements. The mixture being electrolysed is called an electrolyte and must be liquid (either

melted or dissolved) to allow the ions to move.

Electrolysis cell

Electrodes

• General principle that metals or hydrogen are formed at the negative electrode (cathode), and

that non-metals (other than hydrogen) are formed at the positive electrode (anode)

• Describe the electrode products in the electrolysis of:

– molten lead(II) bromide

– concentrated hydrochloric acid

– concentrated aqueous sodium chloride

between inert electrodes (platinum or carbon)

Electrolysis of Molten Ionic Compounds:

An idealized cell for the electrolysis of sodium chloride is shown in the figure below. A source

of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride.

Because the salt has been heated until it melts, the Na+ ions flow toward the negative electrode

and the Cl- ions flow toward the positive electrode.

Negative electrode (cathode): Na+ + e

- → Na

Cl- ions that collide with the positive electrode are oxidized to Cl2gas,

which bubbles off at this electrode.

Positive electrode (anode): 2Cl- → Cl2 + 2e

-

The net effect of passing an electric current through the molten salt in

this cell is to decompose sodium chloride into its elements, sodium

metal and chlorine gas.

2NaCl(l) → 2 Na(l) + Cl2(g)

This example explains why the process is called electrolysis. The suffix -lysis comes from the

Greek stem meaning to loosen or split up. Electrolysis literally uses an electric current to split a

compound into its elements.


Electrolysis of Aqueous Ionic Compounds:

Electrolysing an ionic compound in its solution is very much different to electrolysing it

when it’s molten. This is because in a solution we have 4 ions, H+and OH

- from water

and a positive and a negative ion from the compound. But only one type of ions gets

discharged at each electrode.

For the positive ions, the one that gets discharged at the cathode is the least reactive

one. This is because least reactive elements have more tendencies to be an atom.

So if the ion from the ionic compound is above hydrogen in the reactivity series (more

reactive), H+ gets discharged at the anode And if the ion from the compound is below

hydrogen in the reactivity series (less reactive), this ion gets discharged at the cathode.

So for example if we are electrolysing aqueous sodium chloride, H+ ions will get

discharged at the cathode because sodium is more reactive than hydrogen. And if we

are electrolysing aqueous copper iodide, Cu2+

ions will get discharged at the cathode

because copper is less reactive than hydrogen.

For the negative ions however it is different. Oxygen from OH- from water is always discharged

at the anode except in one case, this is if the other negative ion is a halide. If oxygen from OH-

is discharged, the equation will be:

4OH- - 4e → O2 + H2O

If the other negative ion is a halide, there are two probabilities:

1. Oxygen from OH- gets discharged at the cathode,

2. The halide ion gets discharged at the cathode.

It all depends on the concentration of the halide. If the electrolyte is a concentrated solution,

then there are many of the halide ions, more than OH-. So the halide ion gets discharged at the

cathode. If the electrolyte is a dilute solution, then there are more OH- ions than halide ions, so

oxygen from OH- gets discharged.


So for example if the electrolyte is a concentrated solution of sodium chloride, hydrogen gas is

formed at the cathode because hydrogen is less reactive than sodium. And chlorine gas is

formed at the anode because the solution is concentrated.

If the electrolyte is a dilute solution of silver sulfate, silver is formed at the cathode because it is

less reactive than hydrogen and oxygen gas is formed at the anode.

• Predict the products of electrolysis of a specified halide in concentrated or dilute aqueous

solution

Water is a weak electrolyte: H2O(l) H+

(aq)+OH–(aq)

Discharge series: Cu2+

, H+,Al

3+, Mg

2+,Na

+

I–,Br

–,Cl

–,OH

–,NO3

–,SO4

2–

Anode(+) - anions(–) - lose e– - oxidation

Cathode(–) - cations(+) – gain e– - reduction

Half equation

Aqueous solutions:

2H+

(aq)+ 2e–

H2(g)

4OH–(aq) O2(g)+ 2 H2O (l) + 4e

–

This table shows some common ionic compounds (in solution), and the elements released when

their solutions are electrolysed using inert electrodes, eg carbon electrodes:

Ionic substance Element at - Element at +

Copper chloride, CuCl2 Copper, Cu Chlorine, Cl2

Copper sulfate, CuSO4 Copper, Cu Oxygen, O2

Sodium chloride, NaCl Hydrogen, H2 Chlorine, Cl2

Hydrochloric acid, HCl Hydrogen, H2 Chlorine, Cl2

Sulfuric acid, H2SO4 Hydrogen, H2 Oxygen, O2

Very dilute solutions of halide compounds

If a halide solution is very dilute (eg NaCl), then oxygen will be given off instead of the

halogen. This is because the halide ions are outnumbered by the hydroxide ions from the water.

• The manufacture of chlorine and sodium hydroxide from concentrated aqueous sodium

chloride (brine)

The ions in solution: Na+, H

+, Cl

–,OH

–

Anode(+): 2Cl–(aq) Cl2(g)+2e

–

Cathode(–): 2H+

(aq) +2e–

H2(g)

remain in solution: Na+, OH

–

• Refining of copper(purification by electrolysis)

electrolysis aqueous copper(II) sulfate using copper electrodes

http://www.bbc.co.uk/education/guides/zk96fg8/revision/3#glossary-z8h4mp3

http://www.bbc.co.uk/education/guides/zk96fg8/revision/3#glossary-zgqy87h


Anode(+):Cu(s) Cu2+

(aq) +2e–

Cathode(–):Cu2+

(aq) +2e–

Cu(s)

the electrolyte remains the same deep blue colour

The pure copper rod is connected to the negative terminal of a battery, and the impure rod is

connected to the positive terminal

The pure copper rod has increased in size, while the impure rod has deteriorated, leaving a pool

of anode sludge at the bottom of the beaker

The electrolysis aqueous copper(II) sulfate using carbon electrodes

Anode(+):4OH–

(aq) O2(g)+ 2 H2O (l) + 4e–

Cathode(–):Cu2+

(aq) +2e–

Cu(s)

the electrolyte gradually loses its blue colour

• The electroplating of metals

How it works

The negative electrode should be the object to be electroplated.

The positive electrode should be the metal that you want to coat the object with.

The electrolyte should be a solution of the coating metal, such as its metal nitrate or

sulfate.

Anode(+):Me(s) Mex+

(aq) +xe–

Cathode(–):Mex+

(aq) +xe–

Me(s)

Me: Ag,Au,Cu,Ni,Sn,Cr

• Uses of electroplating: protection from corrosion - appearance

• The manufacture of aluminium from pure aluminium oxide in molten cryolite

The ore crushed and mixed with NaOH

Al2O3(s) + 2 NaOH(aq) 2NaAlO2(aq) + H2O (l)

The impurities are insoluble

The sodium aluminate heated to make up Al2O3

The Al2O3 dissolved in molten cryolite(Na3AlF6) and CaF2 to lower its m. p. and improves the

conductivity

Anode/graphite(+):2O2–

O2(g)+ 4e–(O2 react with C to form CO2)

Cathode/ graphite( (–):Al3+

+3e–

Al

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The overall reaction:

Al2O3 4Al + 3O2

• Conductors

copper : good conductor – ductile – easily purified by electrolysis

steel-cored aluminium in cables

Al: good conductor – low density – resistant to corrosion

steel : additional strength

Thin – thick wires: larger electric current – heat /melt

• Insulators

plastics : flexible – not biodegradable – high electric current :thermosetting

ceramics: high m.p.- not affected by water /air


•Remember that in an electrolyte, it is the ions that move, not the electrons.

• Remember that when a solution of sodium chloride is electrolysed, hydrogen is formed at the

cathode whereas with molten sodium chloride, sodium is formed.

• Make sure that you know the difference in the products at each electrode when dilute and

concentrated aqueous sodium chloride and molten sodium chloride are electrolysed.

• Remember that in electrolysis the electrodes are usually inert (graphite or platinum). If the

anode is not inert, it will react and decrease in size.

•When asked questions about what you observe during electroplating, the answer expected is

what you see happening at each electrode and any changes in the colour of the electrolyte.

•You do not have to learn the diagram of the shell used to extract aluminium but you should be

able to label the different parts. You should also be able to write half equations for the reactions

at the electrodes.

•It is a common mistake to think that the steel core in electricity cables just conducts electricity.

It is also there to strengthen the cables.

• A common mistake is to think that sulphate ions break up during the electrolysis of aqueous

solutions into sulphur dioxide. In fact, oxygen is given off at the positive electrode (from the

electrolysis of the water)

• If the exam paper shows an electrical circuit to test conduction, observations can also include

what can be seen to be happening in the circuit e.g. ‘the bulb lights up’

•It is a common error to muddle cells with electrolysis. In electrolysis an electric current is used

to decompose the electrolyte. In a cell the different reactivity of the electrodes makes an electric

current flow.

• You do not need to remember details about the construction of a fuel cell, but you may be

asked questions based on diagrams and relevant half equations.


6 Chemical energetics

6.1 Energetics of a reaction

• Exothermic and endothermic reactions

• Bond breaking is endothermic and bond forming is exothermic

•ΔH : kJ/mol

For an exothermic reaction, the enthalpy change is always negative.

in exothermic reactions the reactants are higher than the products

For an endothermic reaction, the enthalpy change is always positive.

in endothermic reactions the reactants are lower than the products

•Bond energy

6.2 Production of energy

• Production of heat energy by burning fuels

coal: very polluting – acid rain – global warming

petroleum: less polluting – global warming

natural gas: less polluting – global warming

hydrogen: non polluting – lot of energy – explosive mixture

•Calorimeter

Using the ideas you learn in physics about specific heat capacity, you may have to calculate the

amount of energy released by one mole of a substance.

Heat evolved = m.c.ΔT


Then calculate heat released per mole:

Heat per mole = heat evolved / moles

*ΔT is the temperature rise, m is the mass of the solution in grams which is assumed to equal its

volume in cm3, c is the specific heat capacity of water which is 4.2 J K

-1 g

-1

Fair testing

When comparing different fuels, it is important to carry out a fair test. Several variables should

be kept constant. They include:

the volume of water used

the starting temperature of the water

the temperature increase

the distance of the flame from the calorimeter

• ΔH=ΣΒbroken - ΣΒformed

• Radioactive isotopes, such as 235

U, as a source of energy

• The production of electrical energy from simple cells.

Electrochemical cell Zn/Cu in dilute H2SO4

( –):Zn(s) Zn2+

(aq) + 2e–

(+): 2H+

(aq) + 2e–

H2(g)

The more reactive Me is always the negative electrode

Disadvantanges:

-lose power(reactants used up)

-bulky

-have to be recharged

-harmful

-difficult to dispose of safety

• Fuel cell: hydrogen (is bubbled through negative electrode) react with oxygen(is bubbled

through positive electrode) to generate electricity.

Acidic electrolyte ( H+ produced at the negative electrode and reacts at the positive)

( –):2H2(g) 4H+

(aq) + 4e–

(+):O2(g) +4H+

(aq) + 4e–

2H2O(l)

Alkaline electrolyte ( OH–reacts at the negative electrode and produced at the positive)

( –):2H2(g) +4OH–

(aq) 4H2O(l) + 4e–

(+):O2(g) +2H2O(l) + 4e–

4OH–

(aq)

Advantanges: no pollutants are formed – more energy /gr – lightweight – not recharging – high

efficiency


•Remember that burning is exothermic.

• If asked whether a reaction is endothermic or exothermic, remember the following:

endothermic – heat is put in (e.g. you have to heat with a Bunsen to get a reaction);

exothermic – heat is given out (e.g. burning fuels and neutralisation reactions are always

exothermic)


7 Chemical reactions

7.1 Rate (speed) of reaction

• Measuring rate of reaction

mass of the reaction mixture

volume of gas

amount of light transmitted

change the pH

change in pressure

time taken for a precipitate to make a letter disappear

• Controlled variables – Independent variable

• Calculating rate of reaction

rate of reaction =

Graf: near the start reaction is faster- then gets slower- finally reaction stops

A reaction stops when the limiting reactant is completely used up

• The effect of particle size- surface area on the rate of reactions

Increasing the surface area of a solid reactant increases the rate of reaction

Smaller particles of solid have a larger surface area than larger ones with the same total volume

Danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. mines)

• The effect of catalysts - enzymes on the rate of reactions

A catalyst speeds up the rate of a chemical reaction but is not used up itself

We need tiny amounts of catalyst

There are 2 types of catalyst: solid – in solution

A solid catalyst works by allowing the reactants to get close together

The reaction occurs more quickly at a lower temperature


Activation energy

Activation energy is the minimum energy needed for a reaction to occur when two particles

collide. It can be represented on an energy level diagram.

• The effect of concentration on the rate of reactions: C ↑ rate ↑

Collision theory: Enough energy – number of successful collisions per second

C ↑ frequency of collisions↑ rate ↑

Reactions involving gases : P↑ means C ↑

• The effect of temperature on the rate of reactions: T↑ rate ↑

Collision theory:

T↑ E↑ frequency of collisions↑ rate ↑

Activation energy Ea

T↑ more particles have E>Ea number of effective collisions ↑ rate ↑(more important)

the reactant particles move more quickly

they have more energy

the particles collide more often, and more of the collisions are successful

the rate of reaction increases


• Photochemical reactions ( the effect of light on the rate of reactions)

• The use of silver salts in photography as a process of reduction of silver ions to silver

2Ag+Br

- (crystal) + hv (radiation) 2Ag + Br2

2Ag+ + 2e

- 2Ag : reduction

2Br- Br2 + 2e

- : oxidation

• Photosynthesis the reaction between carbon dioxide and water in the presence of

chlorophyll(catalyst) and sunlight to produce glucose and oxygen

6 CO2 + 6 H2O + Light C6H12O6 + 6 O2


• Many students have difficulty explaining what is meant by rate of reaction. Remember two

points: it is the change in volume or mass etc over a fixed period of time. Time is often omitted

• Remember that the total volume of gas released by the same amount of metal is always the

same. A common error is to think that powdered metal, when reacted with acid, gives off more

gas than larger lumps of the same amount of metal

• The total volume of gas released by a catalysed reaction is exactly the same as for an

uncatalysed reaction. The same amount of reactants is the important factor

• In rate questions, when asked to analyse graphs of volume of gas against time for the reaction

of an acid with a metal or carbonate, a common error is to state the volume is increasing and not

mention the rate. Remember that the rate is getting less and less with time because rate is the

difference in volume divided by time

•Remember that rate of reaction depends on two things: 1. the change in amount or

concentration of reactants or products and 2.the time taken for this change to occur.

•Make sure that you know how to interpret the different parts of a graph of volume of gas

released or loss in mass of the reactants against time. For the Extension you should also be able

to calculate the rate of reaction from these graph.

•It is a common error to think that larger particles have a larger surface area than smaller ones.

Think of a large cube cut up –by cutting, you are exposing more surfaces.

•When defining a catalyst, the best answer is ‘a substance that speeds up a reaction but remains

chemically unchanged at the end of the reaction’. Phrases such as ‘a substance which changes

the rate of a reaction’ are rather vague.

•When explaining the effect of concentration on reaction rate don’t just refer to more collisions

between the particles. It is the more frequent collision of the particles which is important.

• When writing answers to questions about rates of reaction, it is important to use words like

faster or slower not just fast or slow.

•Note that as temperature increases, each particle collides with a greater force. It is also more

accurate to write that there are more frequent collisions than just more collisions.

•It is important to realize that light only affects a few reactions. The only ones you have to

know about are the photosynthesis, the conservation of silver bromine to silver and the reaction

of alkanes with clorine.


8 Reversible reactions

Are ones that can go forward and backwards depending on the conditions

Dehydration and Hydration:

Assume we have a hydrated salt, copper sulphate for example. If you heat the salt you get

two products. They are water and anhydrous copper sulphate. This is a reversible

reaction because if you cool the mixture of the products again, you get hydrated copper

sulphate back.

CuSO4 . 5H2O⇋ CuSO4 + 5H2O

→Heating→

←Cooling←

Note: hydrated copper sulphate is blue crystals. Anhydrous copper sulphate is white powder but

it forms a blue solution with water.

Equilibrium:

Some reversible reactions are very unique, at a certain point, the reaction will be going

forward and backwards at the same time and at the same rate. This is called the state of

equilibrium. In the state of equilibrium, the rate of forward reaction is equal to the rate of

backward reaction and the amount of products and reactants remain constant.

Dynamic equilibrium

1. Rate of forward reaction = rate of reverse reaction

2. Concentrations of all reactants and products remain constant.

3. The system is closed, and on the large scale (macroscopic) everything is constant.

• the effect of changing the concentration, on reversible reactions

Increasing the concentration of a reactant moves the reaction in the direction of the products

If we remove the products from an equilibrium mixture, more reactants are converted into

products.

If a catalyst is used, the reaction reaches equilibrium much sooner, because the catalyst speeds

up the forward and reverse reactions by the same amount.

• the effect of changing the temperature on reversible reactions

If the temperature is increased, the position of equilibrium moves in the direction of the

endothermic reaction

if the temperature is reduced, the position of equilibrium moves in the direction of the

exothermic reaction

• the effect of changing the pressure on other reversible reactions

If the pressure is increased, the position of equilibrium moves in the direction of the

fewest moles of gas.


• Make sure that you understand the term hydrated anhydrous and the water of crystallization

• Remember that if the equilibrium conditions are changed the reaction always tries to act in

the opposite direction

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• A common mistake is to say that in an equilibrium reaction, a catalyst increases the rate of the

forward reaction more than the back reaction. One of the characteristics of equilibrium is that

the backward and forward reactions go at the same speed. This applies to catalyzed as well as

unanalyzed reactions


9 Redox

Oxidation is gain of oxygen.

Reduction is loss of oxygen.

Oxidation is the loss of electrons from a substance. It is also the gain of oxygen by a substance

Reduction is the gain of electrons by a substance. It is also the loss of oxygen from a substance.

Usually, oxidation and reduction take place at the same time in a reaction. We call this type of

reaction a redox reaction.

Note that:

the oxidising agent is the chemical that causes oxidation

the reducing agent causes the other chemical to be reduced

• In a redox reaction involving ions, tow half equations can be writen

• Redox reactions by changes in oxidation state

Assigning oxidation numbers

Rule Examples

1. The oxidation number of each

atom in a pure element is zero.

Zn, O in O2, and P in P4 all have an oxidation

number of zero.

2. The oxidation number of an atom

in a monatomic ion is equal to the

charge on the ion.

Na+ has an oxidation number of +1.

S2-

has an oxidation number of -2.

3. In compounds containing

oxygen, each oxygen atom has an

oxidation number of -2

In H2O and CO2 each oxygen atom has an

oxidation number of -2.

4. In compounds containing

hydrogen, each hydrogen atom has

an oxidation number of +1

In NH3 and H2O each hydrogen atom has an

oxidation number of +1.

5. For a molecule, the sum of the

oxidation numbers of the atoms

equals zero.

The sum of the oxidation numbers of the

atoms in CH4 is zero. As such hydrogen atom

has an oxidation number of +1, the oxidation

number of the carbon atom is -4: (x + (4x + 1)

= 0, x = -4).

6. For a polyatomic ion, the sum of

the oxidation number of the atoms

equals the charge in the ion.

The sum of the oxidation numbers of the

atoms in PO43-

is -3. As each oxygen atom has

an oxidation number of -2, the oxidation

number of the phosphorus atom is +5: (x + (4x

– 2) = -3, x = +5).

7. In a compound, the most

electronegative atom is assigned the

negative oxidation number.

In SF6, the oxidation number of each fluorine

atom is -1. The oxidation number of the sulfur

atom is +6: (x + (6x – 1) = 0, x = +6).

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Oxidation is an increase in oxidation state

Reduction is a reduction in oxidation state

• redox reactions by the colour changes involved when using

acidified potassium manganate(VII) in acidic solution is a good oxidant, when it oxidizes a

substance is color change from purple to colourless

potassium iodide in acidic solution is a good reductant, when it reduces a substance is color

change from colourless to brown


• When explaining redox reactions, make sure you understand exactly what is being asked,

especially if the question says ‘use the equation…’. Don’t just give a definition of redox in

terms of electron loss or gain. If a question says ‘use the equation to explain why the iron oxide

is reduced’, you must refer to the species in the equation in your answer, e.g. ‘the iron oxide

loses its oxygen’. ‘Iron oxide gains electrons’ is incorrect


10 Acids, bases, salts

10.1 The characteristic properties of acids and bases

Acids are substances made of a hydrogen ion and non-metal ions. They have the following

properties:

They dissolve in water producing a hydrogen ion H+,

They have a sour taste,

Strong ones are corrosive,

Their pH is less than 7.

Turns blue litmus paper/ solution red

All acids must be in aqueous form to be called an acid. For example Hydrochloric acid is

hydrogen chloride gas dissolved in water. The most common acids are:

Hydrochloric acid HCl,

Sulphuric Acid H2SO4,

Nitric Acid HNO3,

Cirtric Acid,

Carbonic Acid H2CO3.

Dilute acids react with relatively reactive metals such as magnesium, aluminium, zinc and

iron. The products of the reaction are a salt plus hydrogen gas.

metal + acid → salt + hydrogen

In general, the more reactive the metal, the faster the reaction.

However, aluminium has a protective oxide layer, so it reacts slowly with acids to begin with.

Acids react with metal oxides and hydroxides, a salt and water are made:

acid + metal oxide → salt + water

Acids react with carbonates, such as calcium carbonate (found in chalk, limestone and

marble), a salt, water and carbon dioxide are made. In general:

acid + metal carbonate → salt + water + carbon dioxide

Bases are substances made of hydroxide OH- ions and a metal. Bases can be made of:

Metal hydroxide (metal ion & OH- ion)

Metal oxides

Metal carbonates (metal ion & CO32-

)

Metal hydrogen carbonate (Bicarbonate)

Ammonium hydroxide (NH4OH)

Ammonium Carbonate ((NH4)2CO3)

Properties of bases:

Bitter taste

Soapy feel

Have pH’s above 7

Strong ones are corrosive

Turns red litmus paper/ solution blue


Some bases are water soluble and some bases are water insoluble. Water soluble bases are also

called alkalis.

Reactions of bases

Alkalis react with acids to produce a salt and water (neutralization)

e.g. NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

Metal oxides react with acids to produce a salt and water (neutralization)

e.g. MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)

Metal carbonates react with acids to produce a salt, water and carbon dioxide

e.g. Na2CO3(s) + 2HCl(aq) 2NaCl(aq) + H2O(l) + CO2(g)

Metal hydrogen carbonates react with acids to produce a salt, water and carbon dioxide

e.g. NaHCO3(s) + HCl(aq) NaCl(aq) + H2O(l) + CO2(g)

Displacement of ammonia from ammonium salts

NH4Cl(s) + NaOH(aq) NaCl(aq) + H2O(l) + NH3(g)/(aq)

Ammonia reacts with acids to produce an ammonium salt

e.g. NH3(aq) + HCl(aq) NH4Cl(aq)

• Neutrality and relative acidity and alkalinity in terms of pH - Measured using Universal

Indicator paper

Controlling Soil pH:

If the pH of the soil goes below or above 7, it has to be neutralized using an acid or a base. If

the pH of the soil goes below 7, calcium carbonate (lime stone) or calcium oxide (lime) is

used to neutralize it.

The pH of the soil can be measured by taking a sample from the soil, crushing it, dissolving in

water then measuring the pH of the solution.

Acids and bases in terms of proton transfer

acid is a hydrogen ion (proton) donor.

base is a hydrogen ion (proton) acceptor


Strong acid: an acid that ionizes completely in aqueous solution.

e.g. HCl, HNO3, H2SO4

Weak acid: an acid that ionizes to a small extent (partially)in aqueous solution.

Strong base: a base that almost completely dissociated in aqueous solution , are group 1

hydroxides (ie NaOH etc), or lower group 2 hydroxides Ba(OH)2.

e.g. NaOH, KOH, Ba(OH)2

Weak base: a base that accepts a hydrogen ion from water with difficulty.

Distinguish between equimolar solutions of strong and weak acid.

Strong acid:

has a higher conductivity, better electrical conductor

react more rapidly with magnesium

results to a greater increase in temperature during the neutralization

has lower pH, higher concentration of H+

10.2 Types of oxides

Acidic Oxides

They are all non-metal oxides except non-metal monoxides

They are gases

They react with an alkali to form salt and water

Note: metal monoxides are neutral oxides

Examples: CO2, NO2, SO2 (acidic oxides) & CO, NO,H2O (neutral oxides)

Basic Oxides

They are metal oxides

They react with acids forming a salt and water

They are solids

They are insoluble in water except group 1 metal oxides.

They react with an acid forming salt and water

Examples: Na2O, CaO and CuO

Amphoteric Oxides

These are oxides of Aluminum, Zinc & Lead

They act as an acid when reacting with an alkali & vice versa

Their element’s hydroxides are amphoteric too

They produce salt and water when reacting with an acid or an alkali.

Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l)

Al2O3(s) + 2NaOH 2NaAlO2(aq) + H2O(l)

ZnO(s) + 2HCl(aq) ZnCl2(aq) + 2H2O(l)


ZnO(s) + 2NaOH Na2ZnO2(aq) + H2O(l)

Neutral Oxides

These are N2O, NO,CO

They do not act as an acid or base

10.3 Preparation of salts

Soluble Insoluble

All nitrates None

All common sodium, potassium and ammonium salts None

Most common sulfates Calcium , Barium and Lead

Most common chlorides, bromides, iodides Silver , Lead

Sodium, Potassium and Ammonium Most common carbonates

Group I and Ammonium, (Calcium is slightly soluble) Hydroxides

Group I and Group II react with water Most metal oxides

Preparing Soluble Salts:

Displacement Method (Excess Metal Method):

Metal + Acid → Salt + Hydrogen

Note: this type of method is suitable to for making salts of moderately reactive metals because

highly reactive metals like K, Na and Ca will cause an explosion. This method is used with the

MAZIT (Magnesium, Aluminum, Zinc, Iron and Tin) metals only.

Example: set up an experiment to obtain magnesium chloride salt.

Mg + 2HCl → MgCl2 + H2

Observations of this type of reactions:

Bubbles of colorless gas evolve (hydrogen). To test approach a lighted splint if hydrogen is

present it makes a pop sound

The temperature rises (exothermic reaction)

The metal disappears

You know the reaction is over when:

No more gas evolves

No more magnesium can dissolve

The temperature stops rising

The solution becomes neutral

Proton Donor and Acceptor Theory:

When an acid and a base react, water is formed. The acid gives away an H+ ion and the base

accepts it to form water by bonding it with the OH- ion. A hydrogen ion is also called a proton

this is why an acid can be called Proton Donor and a base can be called Proton Acceptor.


Neutralization Method:

Acis + Base → Salt + Water

Note: This method is used to make salts of metals below hydrogen in the reactivity series. If the

base is a metal oxide or metal hydroxide, the products will be salt and water only. If the base is

a metal carbonate, the products will be salt, water and carbon dioxide.

Type 1: Acid + Metal Oxide → Salt + Water

To obtain copper sulfate salt given copper oxide and sulfuric acid:

CuO + H2SO4 → CuSO4 + H2O

Observations of this reaction:

The amount of copper oxide decreases

The solution changes color from colorless to blue

The temperature rises

You know the reaction is over when

No more copper oxide dissolves


The solution become neutral

Type 2: Acid + Metal Hydroxide → Salt + Water

to obtain sodium chloride crystals given sodium hydroxide and hydrochloric acid:

HCl + NaOH → NaCl + H2O

Observations:

Sodium hydroxide starts disappearing

Temperature rises



No more sodium hydroxide can dissolve

The pH of the solution becomes neutral

Type 3: Acid + Metal Carbonate → Salt + Water + Carbon Dioxide

To obtain copper sulfate salt given copper carbonate and sulfuric acid:

CuCO3 + H2SO4 → CuSO4 + H2O + CO2

Observations:

Bubbles of colorless gas (carbon dioxide) evolve, test by approaching lighted splint, if the

CO2 is present the flame will be put off

Green Copper carbonate starts to disappear

The temperature rises

The solution turns blue


You know the reaction is finished when:

No more bubbles are evolving


No more copper carbonate can dissolve

The pH of the solution becomes neutral

Titration Method:

This is a method to make a neutralization reaction between a base and an acid producing a salt

without any excess.

is used to make a soluble salt

the experiment is preformed twice, the first time, using an indicator ,is to find the amounts of

reactants to use, and the second experiment is the actual one.

Other indicators

Indicator Acidic Neutral Alkaline

Methyl orange Red Yellow Yellow

Phenolphthalein Colourless Colourless Pink

Preparing Insoluble Salts:

Precipitation Method:A precipitation reaction is a reaction between two soluble salts. The

products of a precipitation reaction are two other salts, one of them is soluble and one is

insoluble (precipitate).

Example: To obtain barium sulfate salt given barium chloride and sodium sulfate:

BaCl2 + Na2SO4 → BaSO4 + 2NaCl

Ionic Equation: Ba2+

+ SO42-

→ BaSO4

Observations:

Temperature increases

An insoluble solid precipitate (Barium sulfate) forms




No more precipitate is being formed

• Suggest a method of making a given salt from suitable starting material, given appropriate

information

10.4 Identification of ions and gases

Colors of Salts:

Salt Formula Solid In Solution

Hydrated copper

sulfate

CuSO4.5H2O Blue crystals Blue

Anhydrous copper

sulfate

CuSO4 White powder Blue

Copper nitrate Cu(NO3)2 Blue crystals Blue

Copper chloride CuCl2 Green Green

Copper carbonate CuCO3 Green Insoluble

Copper oxide CuO Black Insoluble

Iron(II) salts E.g.: FeSO4, Fe(NO3)2 Pale green

crystals

Pale green

Iron(III) salts E.g.: Fe(NO3)3 Reddish brown Reddish brown

Tests for Gases:

Gas Formula Tests

Ammonia NH3 Turns damp red litmus paper blue

Carbon

dioxide

CO2 Turns limewater milky

Oxygen O2 Relights a glowing splint

Hydrogen H2 ‘Pops’ with a lighted splint

Chlorine Cl2 Bleaches damp litmus paper

Nitrogen

dioxide

NO2 Turns damp blue litmus paper red

Sulfur dioxide SO2

Turns acidified aqueous potassium dichromate(VI) from

orange to green

Tests for Anions:

Anion Test Result

Carbonate (CO32-

) Add dilute acid Effervescence,

carbon dioxide produced

Chloride (Cl-)(in

solution)

Acidify with dilute nitric acid, then

add aqueous silver nitrate

White ppt.

Iodide (I-)(in solution) Acidify with dilute nitric acid, then

add aqueous silver nitrate

Yellow ppt.

Nitrate (NO3-)(in

solution)

Add aqueous sodium hydroxide,

then aluminium foil; warm

carefully

Ammonia produced

Sulfate (SO42-

) Acidify, then add aqueous barium

nitrate

White ppt.


Tests for aqueous cations:

Cation Effect of aqueous sodium

hydroxide

Effect of aqueous ammonia

Aluminium (Al3+

) White ppt., soluble in excess giving

a colourless solution

White ppt., insoluble in excess

Ammonium

(NH4+)

Ammonia produced on warming –

Calcium (Ca2+

) White ppt., insoluble in excess No ppt. or very slight white ppt.

Copper (Cu2+

) Light blue ppt., insoluble in excess Light blue ppt., soluble in excess,

giving a dark blue solution

Iron(II) (Fe2+

) Green ppt., insoluble in excess Green ppt., insoluble in excess

Iron(III) (Fe3+

) Red-brown ppt., insoluble in excess Red-brown ppt., insoluble in excess

Zinc (Zn2+

) White ppt., soluble in excess,

giving a colourless solution

White ppt., soluble in excess,

giving a colourless solution



• Don’t confuse the pH scale with the degree of acidity. The more acidic the substance, the

lower the pH – learn this by remembering that ‘a’ (for acid) is the lowest numbered letter of the

alphabet

• A common error is to think that less sodium hydroxide is needed to neutralise a weak acid

than to neutralise a strong acid of the same concentration. The same amount is needed because

the hydroxide is reacting with all the acidic hydrogens in the molecule, not just those that have

ionised

• The phrase ‘explain why this acid is acting as a base’ demands a chemical reason (usually

based on particle theory). The examiner is looking for an answer involving proton transfer.

Vague answers (such as ‘it is neutralising the base’) are not accepted as they do not give an

explanation

• Simple inorganic salts such as sodium chloride are generally neutral when dissolved in water

– they are not acidic

• Nitric acid is a strong, not a weak, acid

• A common error is to think that calcium hydroxide is insoluble in water. Remember that

limewater is a solution of calcium hydroxide, so it must at least be slightly soluble

• If you are asked to explain what the symbol (aq) means, write down more than ‘aqueous’. An

answer such as ‘dissolved in water’ is needed

• Look out for phrases such as ‘chemical test’ or ‘physical test’ – don’t just focus on the word

‘test’. For example, a chemical test for water could be ‘turns anhydrous copper sulphate blue’

(the word ‘anhydrous’ is essential). A physical test for water could be ‘a boiling point of

100oC’, using the correct units


• When testing hydrogen chloride gas with litmus paper, many students think that the litmus

paper is bleached first and then goes red. Remember that chlorine does this, not hydrogen

chloride

• The tests for ammonium and nitrate ions are commonly confused. Both require heating with

sodium hydroxide, but to test for nitrate you need to add aluminium, as you need to remove the

oxygen (reduce the nitrate) to make the ammonia. You don’t need to do this for the ammonium

ion as it has no oxygen

• Tests for aluminium ions and zinc ions are also often confused. Remember PANDA

(precipitate of aluminium (hydroxide) does not dissolve in ammonia). Both zinc and aluminium

ions form a white precipitate with sodium hydroxide, which re-dissolves in excess, but in

ammonia only the zinc precipitate re-dissolves

• Questions involving the height of precipitates when sodium hydroxide is added to a solution

of metal ions often cause problems. Remember, as you add more hydroxide to a solution of

suitable metal ions (e.g. iron(II) ions) there will be more precipitate until all the metal ions are

used up. However, with excess sodium hydroxide, some hydroxides re-dissolve e.g. aluminium

hydroxide. In these cases the height of the precipitate will then decrease as you add more

hydroxide

•Remember that a lower acidity gives a higher pH and a higher acidity gives a low pH.

•Don’t forget that when acids react with carbonates, water is produced – as well as a salt and

carbon dioxide. For extension you must be able to write the symbol equations.

•It is incorrect to use the word ‘strong’ and ‘weak’ when referring to the concentration of acid

or alkalis. Use ‘concentrated’ or ‘dilute’. Strong and weak refer to the degree of ionization of

the acid or base, not the concentration.

•When you make a salt using excess metal or metal oxide, you first have to filter off the excess

solid reactant. You may be asked how to make a salt in any of the exam papers.

•Make sure that you know what types of compound are soluble or insoluble. Without this

knowledge you will not be able to select precipitation as the correct method to make a particular

salt.

•A common error is to confuse the tests for hydrogen and oxygen. It may help you to remember

that ‘lighted’ (splint) has an ‘h’ in it for hydrogen and ‘glowing’ (splint) has an ‘o’ in it for

oxygen.

•When testing for metal ions using sodium hydroxide, make sure that you mention three things:

(i) if there is a precipitate (ii) the colour of the precipitate (iii) what happens when you add

excess sodium hydroxide

•Remember that you add nitric acid and silver nitrate in the test for halide ions. If you add

hydrochloric acid you will be adding chloride ions!


11 Periodic table

• The Periodic Table as a method of classifying elements and its use to predict properties of

elements

Periodic trends

• the change from metallic to non-metallic character across a period

• the relationship between Group number, number of valency electrons and metallic/non-

metallic character

Special Elements:

Alkali Metals:

These elements lie in group 1 of the periodic table. They are Lithium, Sodium, Potassium,

Rubidium, Caesium and Francium (radioactive). We will study the properties of the first three;

Lithium, Sodium and Potassium. Like any metals they are all good conductors of heat and

electricity. They are however, soft. Lithium is the hardest of them and potassium is the softest.

They are extremely reactive; they have to be stored away from any air or water. They have low

densities and melting points.

They react with oxygen or air forming a metal oxide:

4Li +O2 → 2Li2O

Their oxides can dissolve in water forming an alkaline solution of the metal hydroxide:

Li2O + H2O → 2LiOH

(Lithium Oxide) (Water) (Lithium Hydroxide)

They react with water vigorously forming metal hydroxide and hydrogen gas:

2K + 2H2O → 2KOH + H2

They React with Halogens forming a metal halide:

2Na + Cl2 → 2NaCl

The reactivity of Group 1 elements increases as you go down the group because:

the atoms get larger as you go down the group

the outer electron gets further from the nucleus as you go down the group

the attraction between the nucleus and outer electron gets weaker as you go down the group

- so the electron is more easily lost


Flame colors and the alkali metal ion they represent

Flame colour Ion present

Red Lithium, Li+

Orange Sodium, Na+

Lilac Potassium , K+

Brick red Calcium, Ca2+

The Halogens:

These are elements of group 7; Fluorine, Chlorine, Bromine, Iodine and Astatine.

We will study only properties of chlorine, bromine & iodine. They are colored and the color

gets darker as we go down the group. They exist as diatomic molecules (Cl2, Br2, I2). As you go

down, they gradually change from gas to solid (chlorine is gas, bromine is liquid and iodine is

solid).

They react with hydrogen forming hydrogen halide, which is an acid if dissolved in water:

H2 + Cl2 → 2HCl

(Hydrogen) (Chlorine) (Hydrochloric Acid)

They react with metals forming metal halide:

2Fe + 3Cl2 → 2FeCl3

The reactivity also decreases as we do down, chlorine is most reactive, followed by bromine

then iodine.

If you bubble chlorine gas through a solution of potassium bromide, chlorine will take

bromine’s place because it more reactive. This is a displacement reaction.

2KBr + Cl2 → 2KCl + Br2

Transition Elements:

These are metals. They form a big part of the periodic table. Some of them are very common

like copper, zinc and iron. They have the following properties:

They are harder and stronger than metals of groups 1 & 2.

They have much higher densities than metals other metals.

They have high melting points except for mercury.

They are less reactive than metals of group 1 & 2.

Excellent conductors of heat and electricity.

They show catalytic activity (act as catalysts)

They react slowly with oxygen and water

They form simple ions with several oxidation states and complicated ions with high oxidation

states.

They form coloured compounds


Noble Gases:

These are elements in group 8 of the periodic table.

They are colorless gases.

They are extremely unreactive; this is because they have their outer energy shell full with

electrons. So they are stable, this is why they exist as single atoms.

Noble

gas Uses

Helium

Party balloons, airships, cooling superconducting electromagnets (eg in

MRI scanners), gas for scuba diving

Neon Red neon signs, lasers

Argon

Shielding gas for welding, surrounding the filament in an old-fashioned

lightbulb

Xenon Lights, lasers

Krypton Lights, photographic flashguns


12 Metals

12.1 Properties of metals

Metallic bonds

Properties of Metals

Good conductors of electricity : metals have delocalized electrons / sea of electrons

which are mobile.

Good conductors of heat : electrons jumping through cations and moving energy

from M to M .

Shiny : light absorbed by electrons and re-emitted at different Energy levels.

Malleable : pushing layers, atoms/ions/layers (of positive ions) can slide over each

other without change in the bonding forces /

Ductile : moving the layers.

(impurities – alloys: harder than the pure metals)

Melting and boiling point : high

• Alloys

An alloy is a mixture of two or more elements, where at least one element is a metal. Most

alloys are mixtures of two or more metals

Alloys contain atoms of different sizes. These different sizes distort the regular arrangements of

atoms. This makes it more difficult for the layers to slide over each other, so alloys are harder

than the pure metal.

It is more difficult for layers of atoms to slide over each other in alloys

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12.2 Reactivity series

Reactions with Dilute Hydrochloric Acid:

Metal + HCl → Metal Chloride + Hydrogen

Metals Reactivity with Dilute HCl

Potassium, Sodium &

Calcium

React extremely violently with rapid effervescence and

splashing

Magnesium & Aluminum React violently with rapid effervescence

Zinc, Iron & Lead React slowly with bubbles

Copper, Silver, Gold &

Platinum

Do not react

Reactions with Oxygen in Air:

Most metals react with oxygen from air forming a metal oxide.

Metal React with oxygen Product

potassium, sodium, calcium

and magnesium

with a very bright

flame

white ashes and their oxides are

soluble.

aluminum and zinc white powdered ashes but their oxides

are insoluble. A layer of aluminum

oxide adheres and covers the

aluminum. At this point no further

reaction can take place.

Iron and copper very slowly rust which is reddish brown iron oxide

- insoluble

copper lump a white layer of black copper oxide

forms on it. When the lump gets

covered by this layer; the reaction

stops- insoluble

silver, gold and platinum do not react

Reactions of Metals with Water and Steam:

Potassium, sodium and calcium react vigorously with cold water and may catch on fire. The

products of these reactions are metal hydroxide and hydrogen gas. If hydrogen gas being

produced accumulates it may ignite and cause an explosion.

Metal + Water → Metal hydroxide + Hydrogen

E.g.: 2Na + 2H2O → 2NaOH + H2

Magnesium, aluminum, zinc and iron are less reactive. They react with steam forming metal

oxide and hydrogen. Magnesium and aluminum will react vigorously with steam while zinc and

iron react slowly.

Metal + Steam → Metal Oxide + Hydrogen

E.g.: Magnesium + Steam → Magnesium oxide + Hydrogen

Unreactive metals such as silver and gold do not react with water.


Single Displacement Reactions in Solid State:

Previously you’ve studied displacement reactions which are pre-formed in aqueous states. A

very similar reaction takes place in the solid state, it is called thermite reaction. This reaction

is used to repair damaged railway lines. In this reaction, aluminum and iron (III) oxide are the

reactants. In the reaction, aluminum removes the oxygen ion from iron and bonds with it. This

happens because aluminum is more reactive than iron. The products are aluminum oxide and

iron in molten form. In the fixing procedure, the reactants are put in the cut in the railway line

and the reaction is triggered by heating using a magnesium fuse. The reaction leaves aluminum

oxide and molten iron with then condenses in the cut welding it. Like displacement reactions,

this reaction is exothermic.

2Al + Fe2O3 → Al2O3 +2Fe

Single Displacement Reactions in Aqueous State:

These are ordinary displacement reactions in which the two positive ions compete for the

negative ion. The ion of the more reactive metal wins. Zinc is higher than copper in the

reactivity series. If zinc is added to a solution of copper nitrate, a displacement reaction will

take place in which the zinc will displace the copper ion from the solution in its salt. The

products of this reaction are zinc nitrate and copper. Copper salt solutions have a blue color

which fades away as the reaction proceeds because the concentration of the copper salt

decreases. This type of reaction also helped in confirming reactivity of metals since the more

reactive metal displaces the less reactive one.

Zn + Cu(NO3)2 → Zn(NO3)2 + Cu

Explaining reactivity

The easy with which a metal loses its valency electrons depends on the

distance of the valency electrons from the nucleus

the nuclear charge (number of protons)

the number of electrons shells

Reducing metal oxides with carbon

Metal oxides below C in the reactivity series are reduced by carbon when heated

Action of Heat on Metal Compounds:

Applying heat to a metal compound such as potassium nitrate will cause it to decompose into

potassium nitrite and oxygen. This is a thermal decomposition reaction.

Metal: Anion:

Nitrate (NO3) Carbonate (CO3) Hydroxide (OH)

Potassium

Sodium

Metal Nitrate →

Metal nitrite + Oxygen

NO DECOMPOSITION

Calcium

Magnesium

Aluminum

Zinc

Iron

Lead

Copper

Metal Nitrate →

Metal oxide + Nitrogen

dioxide + Oxygen

Metal Carbonate →

Metal oxide + Carbon

dioxide

Metal hydroxide →

Metal oxide +

Hydrogen

Silver

Gold

Metal Nitrate →

Metal + Nitrogen

dioxide + Oxygen

Metal Carbonate →

Metal + Carbon

dioxide + Oxygen

Silver and gold

hydroxides do not

exist.


Ions of more reactive metals tend to hold on tightly to their anions and do not decompose easily

this is why lots of heat is needed.

12.3 Extraction and Uses of metals

Metal Method

Potassium Electrolysis

Sodium Electrolysis

Calcium Electrolysis

Magnesium Electrolysis

Aluminium Electrolysis

(Carbon) (Non-metal)

Zinc Reduction by carbon or carbon monoxide

Iron Reduction by carbon or carbon monoxide

Tin Reduction by carbon or carbon monoxide

Lead Reduction by carbon or carbon monoxide

(Hydrogen) (Non-metal)

Copper Various chemical reactions

Silver Various chemical reactions

Gold Various chemical reactions

Platinum Various chemical reactions


Extraction of Iron:

The ore of iron is called hematite. It consists of 60% iron in form of Iron oxide (Fe2O3) with

other impurities such as silicon oxide (SiO2). This process takes place in a tower called a Blast

furnace.

Substances Products and Waste

Materials

Iron ore (Hematite)

Coke (heated coal)

Lime stone (Calcium

carbonate)

Hot Air

Pure Iron

Carbon dioxide

Air

Slag (Calcium silicate)

Substances are put in the blast furnace

The process starts by blowing in hot air at the bottom of the furnace

Coke burns in oxygen from the hot air producing carbon dioxide; C + O2 → CO2

Heat makes lime stone decompose into calcium oxide and carbon dioxide; CaCO3 → CaO +

CO2

Carbon dioxide produced goes up the furnace and reacts with more coke up there producing

carbon monoxide; CO2 + C → 2CO

Carbon monoxide is a reducing agent. It rises further up the furnace where it meets iron oxide

and starts reducing it producing iron and carbon dioxide; Fe2O3 + 3CO → 2Fe + 3CO2

Calcium oxide which was produced from the thermal decomposition of lime stone is a base. It

reacts with impurities of hematite such as silicon oxide which is acidic forming calcium silicate

which is called slag; CaO + SiO2 → CaSiO3

Molten Iron and slag produced trickles down and settles at the bottom of the furnace. Iron is

denser than slag so it settles beneath it.

Iron and slag are tapped off separately at regular intervals and pure iron is collected alone

Waste gases such as carbon dioxide formed in the process and nitrogen and other gases from

air blown in escape at the top of the furnace.

Conversion of Iron into Steel:

Iron produced in the blast furnace is called pig iron. It contains 4% carbon as well as other

impurities such as sulfur, silicon and phosphorus which make it hard and brittle. It got that

name from the fact that it has to be poured into mould called pigs before it is converted into

steel. Most of produced iron is converted into steel because steel has better properties.

If all the impurities are removed, the iron becomes very soft In this condition, it easily shaped

but is too soft for many uses. Pure iron also rust very easily.

Making steel out of pig iron is a process done in a basic oxygen furnace:

Molten pig iron is poured into the oxygen furnace

A water cooled lance is introduced which blows oxygen onto the surface of the molten iron

Impurities start to react

Carbon is oxidized into carbon monoxide and carbon dioxide and escape

Sulfur is oxidized into sulfur dioxide and escapes

Silicon and phosphorus are oxidized into silicon oxide and phosphorus pentoxide which are

solids.

Calcium oxide (lime) is added to remove the solid impurities as slag which is skimmed off the

surface

Throughout the process, sample of the iron are being taken and analyzed for the percentage of

carbon present in it. When the percentage of carbon desired is reached, the furnace is switched

off and the steel is collected.


There are many different forms of steel.

Steel Composition Properties Uses

Mild Steel 99.5% Iron

0.5% Carbon

Easily worked lost brittleness Car bodies

large structures

Machinery

Hard Steel

99% Iron

1% Carbon

Tough and brittle Cutting tools and

chisels

Stainless Steel

87% Iron

13% Manganese

Tough and springy

Drill bits and springs

and chemical plants

Manganese

Steel

74% Iron

18% Chromium

8% Nickel

Tough and resistant to corrosion Cutlery and surgical

tools, kitchen sinks

Tungsten Steel 95% Iron

5% Tungsten

Tough and hard even at high

temperatures

Edges of high

speed cutting tools

Extraction of Zinc:

The ore of zinc is called zinc blende and it is made of zinc sulfide. Zinc is obtained from zinc

sulfide by converting it into zinc oxide then reducing it using coke, but first zinc sulfide must be

concentrated.

Zinc sulfide from zinc blende is concentrated by a process called froth floatation. In this

process, the ore is crushed and put into tanks of water containing a frothing agent which makes

the mixture froth up. Hot air is blown in and froth starts to form. Rock impurities in the ore get

soaked and sink to the bottom of the tank. Zinc sulfide particles cannot be soaked by water;

they are lifted by the bubbles of air up with the froth and are then skimmed off. This is now

concentrated zinc sulfide.

Then, zinc sulfide gets heated very strongly with hot air in a furnace. Zinc sulfide reacts with

oxygen from the air to produce zinc oxide and sulfur dioxide gas which escapes as waste gas.

2ZnS + 3O2 → 2ZnO + 2SO2

Sulfur dioxide is used in the manufacture of sulfuric acid.

Zinc oxide produced is put into a furnace with powdered coke. The mixture is heated till

1400oC. Carbon from the coke reduces the zinc oxide into zinc producing carbon monoxide

which escapes as waste gas.

ZnO + C → Zn + CO


Carbon monoxide produced is hot and is used to heat the furnace to reduce heating costs. The

pure zinc produced is collected and left to cool down. Zinc is used in many ways like the

production of the alloy brass, galvanization and making car batteries.

Uses of Zinc:

for galvanising and for making brass

Extraction of Aluminum:

Aluminum exists naturally as aluminum oxide (alumina) in its ore, which is called bauxite.

Because aluminum is a very reactive metal, it holds on very tightly to the anion it bonds with,

which is oxide in this case. This is why the best way to extract and purify aluminum is by

electrolysis in a cell like the one below.

In this cell, the electrodes are made of graphite (Carbon). The cathode is a layer at the bottom of

the cell and the anodes are bars dipped in the electrolyte. The electrolyte in this process is a

molten mixture of aluminum oxide and cryolite. Aluminum oxide by its self has a very high

melting point of 2050oC which is higher than the melting point of the steel container in which

this process is done. That means the steel container will melt before the aluminum oxide. This

is why aluminum oxide is mixed with cryolite which decreases the melting point of it to under

1000oC, thus saving a lot of money because heating is expensive and preventing the steel

container from melting. Heat must be continuously supplied to the mixture to keep it molten.

Aluminum oxide does not conduct electricity when solid because it does not have free mobile

ions to carry the charge.

Aluminum oxide is purified from impurities of oxide by adding sodium hydroxide

Aluminum oxide is mixed with cryolite and put in the electrolysis cell

Heat is given in until the mixture becomes molten

Electrolysis start

Oxide ions get attracted to the anode and discharged (oxidation); 2O2-

, 4e → O2

Aluminum ions get attracted to the cathode and discharged and settle at the bottom

of the container (reduction); Al3+

+ 3e → Al

Oxygen gas evolves and is collected with waste gases

Aluminum is sucked out of the container at regular intervals

Oxygen gas which evolves reacts with carbon from the cathode forming CO2. The cathode gets

worn away. To solve this, the cathode is replaced at regular intervals. Heat supply is very

expensive; this is why cryolite is used to decrease the melting point of aluminum oxide and this

process is done in plants which use hydroelectric energy because it is cheap.


Uses of aluminum:

Construction of air-craft bodies because aluminum is very strong and very light and it is

resistant to corrosion

Food containers because it is resistant to corrosion

Overhead power cables because it conducts electricity, is very light, malleable and ductile.

Although it is strengthened with steel core

Extraction of Copper:

Copper is one of the most popular metals. Native copper occurs in some regions in the world.

Otherwise, copper exists in its ore, copper pyrites (2CuFeS2). You have studied before that

copper can be purified by electrolysis. It can also be extracted from it ore by converting pyrites

into copper sulfide by reacting it with oxygen:

2CuFeS2 + 4O2 → Cu2S + 3SO2 + 2FeO

Sulfur oxide produced escapes as waste gas and iron oxide impurities are removed by heating

the mixture with silicon converting it in to iron silicate which is run off. The remaining copper

sulfide is then heated strongly with air. Copper sulfide reacts with oxygen from air producing

sulfur oxide which escapes as waste gas and pure copper.

Cu2S + O2 → 2Cu + SO2

Thus copper is extracted.

Uses of Copper:

In electrical wires because it is a perfect electrical conductor and very ductile, malleable and

cheap

Making alloys such as bronze and brass

Cooking utensils because it conducts heat and it is has high melting and boiling points and

also resists corrosion

Electrodes because it is a good conductor of electricity

Water pipes because it is resistant to corrosion


• Don’t confuse the properties of elements with those of their compounds (especially when they

appear in the same question). For example, if asked about the properties of the element oxygen,

don’t give the properties of an oxide

• The properties of transition elements often cause problems. Remember that transition

elements themselves are NOT coloured, it is their compounds that are coloured

• When trying to distinguish between a transition metal and a non-transition metal, information

on boiling points is more important than information on density. Some non-transition elements

(such as lead) are very dense

• If asked about the specific properties of transition metals, don’t list general properties of

metals, such as ‘shiny’, ‘malleable’, etc.

• In questions about sacrificial protection, remember that the more reactive metal of the pair

will corrode. To answer this sort of question, know the order of common metals in the reactivity

series


• ‘Corrosive’ and ‘corrosion’ are often confused. ‘Corrosive’ means that a chemical ‘eats away’

another substance – acids and alkalis are corrosive. ‘Corrosion’ is the process of ‘eating away’.

A statement such as ‘iron is corrosive’ is therefore incorrect

• The source of an element is where it is found (i.e. a particular place or in a particular

substance) – a source of sulphur is the southern USA, or petrol. It does not mean the process of

extraction. Don’t write vague statements such as ‘underground’

• Sulphur dioxide is not used ‘to make wood pulp’, it is used to bleach wood pulp

•You need to know where the metals and the non-metals appear in the Periodic Table. You do

not have to remember exactly the dividing line between metals and non-metals in Groups III to

VI.

•When you describe observations concentrate on what you see, hear smell or feel by touch.

•When you compare Group I metals, remember that they have ‘similar properties’ NOT ‘the

same properties’. The properties change slightly down the group.

•Make sure that you can distinguish between the halogens (elements) and halides (compounds).

It is a common error to write chlorine ions instead of chloride ions.

•It is better to write that the noble gases are unreactive ‘because they have a full outer shell of

electrons’, which is inaccurate.

•It is a common error to suggest that transition elements are highly coloured. It is the

compounds of transition elements which have a range of colours.

•Remember that oxidation state does not always refer to the charge on the ions. For example, in

potassium manganate (VII), KMnO4 , the oxidation state of manganese is +7 but the manganese

ion with the highest charge is Mn+2.

• It is a common error to think that all metals are hard and have very high melting points.

Remember that Group I metals are soft and have low melting points.

•Remember that metals that react with cold water form metal hydroxides. When a metal is

heated in steam, an oxide is formed.

•In your exam you will usually be given the reactivity series to help you answer questions about

the ease of formation of ions.

•Remember that aluminium is a reactive metal. It must be reactive if it forms an unreactive

oxide layer on its surface so quickly.

•You need to remember the products from the thermal decomposition of nitrates. If you don’t

know these, you won’t be able to write equations for thermal decomposition.

•You will not be asked to draw the furnace used for the extraction of zinc but you should be

prepared to label a diagram and write relevant equations.

• You will not be asked to draw the blast furnace. You should be prepared to answer questions

related to a diagram of the blast furnace and the reactions involved.

•Do not confuse steelmaking with the blast furnace. In steelmaking the impurities are removed

from the impure iron we get from the blast furnace. In the blast furnace the impure iron is

extracted from the iron ore.


13 Air and water

Chemical tests for water

Pure copper(II) sulfate is white. It is also known as anhydrous copper(II) sulfate because it

has no water in it.

When water is present in a sample of copper(II) sulfate it turns blue. It is still a dry solid,

because the individual water molecules are trapped within the ionic lattice surrounding the

copper(II) ions.

Solutions of copper(II) sulfate are also blue.

Water can also be detected using blue anhydrous cobalt(II) chloride. This turns pink in the

presence of water.

Uses of Water:

The uses of water are many, from drinking and cleaning to irrigating crops and landscapes.

Water is used for cooling, for recreation, and dust control. Water is needed for restaurants, most

industrial processes, and even some religious ceremonies. On another level, the splash and flow

of water in streams and fountains soothes and inspires.

In one way or another, water is a part of almost everything humans make and do. Washing a

load of laundry uses 40 gallons, filling a backyard pool takes about 25,000 gallons, growing a

pound of cotton consumes 1,000 gallons, while producing a pound of copper uses 20 gallons.

Uses where water is consumed, usually through evaporation or plant growth, are consumptive

uses. Examples include water used for irrigation or in evaporative coolers. Non-consumptive

uses, such as bathing, hydropower generation and recreation, do no t use up water. Used non-

consumptively, the same water can be used again and again, although some uses lower the

quality of the water. Once used, wastewater can be treated and used again as reclaimed water or

effluent.

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The main categories of water use are agricultural, municipal and industrial. Municipal and

industrial uses currently are much less, but are growing rapidly. Mining activities and cooling

towers used for power generation account for most of the remaining water use.

Water Purification:

1. Water that exists naturally in earth is never pure. There are always impurities in it, sometimes

in large amounts. In fact water could very well be contaminated with diseases and bacteria. This

is why water has to be purified before it is put to use. Water purification involves two processes

(Filtration & Chlorination) done in several steps:

2. Water is taken from reservoirs or any other source to the water treatment plant

3. Water is passed through filters to remove large, floating objects such as pieces of rocks or

mud

4. Smaller particles are removed by adding aluminum sulfate which makes them stick together

in large pieces and settle down

5. Water is passed through sand and gravel filters which filter off small particles and may kill

some bacteria (filtration is done)

6. Chlorine gas is bubbled through the water to kill all bacteria living in the water making the

water sterile

7. The water may end to be slightly acidic, small amounts of sodium hydroxide are added to

treat this. Fluoride might be added to because it helps in preventing tooth decay

8. Water is then delivered to homes

Composition of clean air

Fractional distillation of air

About 78 per cent of the air is nitrogen and 21 per cent is oxygen. These two gases can be

separated by fractional distillation of liquid air.

Liquefying the air

Air is filtered to remove dust, and then cooled in stages until it reaches –200°C. At this

temperature it is a liquid. The air has been liquefied.

Here's what happens as the air liquefies:

water vapour condenses, and is removed using absorbent filters

carbon dioxide freezes at -79°C, and is removed

oxygen liquefies at -183°C

nitrogen liquefies at -196°C

The liquid nitrogen and oxygen are then separated by fractional distillation.

The liquefied air is passed into the bottom of a fractionating column. Just as in the columns

used to separate oil fractions, the column is warmer at the bottom than it is at the top.

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Fractional distillation

Air Pollution:

Pollution is the presence of harmful substances. Air pollution is the presence of pollutant gases

in the air. A pollutant is a substance that causes pollution. These are:

Carbon monoxide

Oxides of nitrogen

Sulphur dioxide

Lead compounds

Carbon Monoxide: Carbon monoxide (CO) is one of the poisonous pollutants of air. It is

considered a pollutant because it can kill living organisms. The main source of carbon

monoxide is factories which burn carbon-containing fossil fuels since CO is one of the products

of the incomplete combustion of fossil fuels. Carbon monoxide could be treated by installing

catalytic converters in chimneys of the factories.

Sulphur Dioxide: Sulphur dioxide (SO2) is considered a pollutant since it contributes to acidic

rain. Sulphur dioxide is a product of two process, these are combustion of sulphur –containing

fossil fuels and extraction of metals from their sulphide ores (such as zinc sulphide). The

problem associated with sulphur dioxide is that when it rises in the air from chimneys of

factories, it mixes with water vapour of clouds and air. This results in the formation of sulphuric

acid (H2SO4). When it rains, rain water which falls becomes acidic. Acid rain causes death to

water creatures since it makes water acidic, acidifies soil causing death to plants and

deforestation, reacting with limestone from buildings and sculptures corroding it, and may also

cause lung cancer. Sulphur dioxide could be treated before it leaves chimneys of factories by

reacting it with limestone which is a neutralisation reaction. This process is called

desulphurisation.

SO2 + CaCO3 → CaSO3 + CO2

Oxides of Nitrogen (NO & NO2): Nitrogen oxides are formed at high temperatures as a result

of nitrogen and oxygen reacting. In cars, engines have a very high temperature; this creates a

chance for nitrogen and oxygen present in air in the engine to react forming nitrogen monoxide.

N2 + O2 → 2NO


The produced carbon monoxide is released through the exhaust with other waste fumes.

Nitrogen monoxide reacts with more oxygen from air producing nitrogen dioxide.

2NO + O2 → 2NO2

The problem associated with nitrogen dioxide is similar to that of sulphur dioxide. It rises up in

the air and mixes with rain water forming nitric acid. This causes acid rain. Nitrogen oxides can

also cause health respiratory problems to humans and animals. To treat this issue, cars are now

fitted with devices called catalytic converters which eliminate nitrogen oxides.

Lead Compounds: Compounds of lead are waste products of fuel burning in cars. They are

considered pollutants because they are poisonous and they are said to cause mental disabilities

to young children. To treat this problem, gas stations now provide unleaded fuel.

Catalytic Converters:

Car fuels contain carbon; so carbon monoxide gas is released by cars as waste fumes, as well as

nitrogen oxides. These are pollutant gases. To prevent these gases from polluting air, a device

called catalytic converter is fitted at the end of the exhaust. This device contains a catalyst

which catalyses the reaction between these two gases producing two harmless gases, nitrogen

and carbon dioxide:

2NO + 2CO → 2CO2 + N2

2NO2 + 4CO → 4CO2 + N2

The catalyst of the device works best at temperature around 200°C.

• State the adverse effect of common pollutants on buildings and on health

The carbon cycle

The carbon cycle is a natural global cycle of the element carbon. It is what maintains a constant

level of carbon dioxide in air (0.03%). The cycle goes as follows:

Plants absorb carbon dioxide from air and undergo photosynthesis reaction which turns it into

glucose and produces

oxygen: 6CO2 + 6H2O → C6H12O6 + 6O2

The carbon is now stored in plants as glucose. One of two things happen, either the plants get

eaten by animals or humans, or the plant dies and decays.

If the plant is eaten by animals or humans, glucose in the plant is used by them in a process

called respiration to release energy for their body.


C6H12O6 + 6O2 → 6CO2 + 6H2O Respiration is the opposite of photosynthesis. Carbon dioxide is one of the products of it, which

is released by the humans through breathing into the air. Thus carbon dioxide returns to the

atmosphere.

If the plant dies. It is buried underground and by time it decays forming coal and other fossil

fuels. These substances contain the carbon which was made and stored by the plants and they

are then taken by power stations which put them to use.

Power stations burn carbon-containing fuels that were obtained as coal or fossil fuels formed by

dead plants. This is a combustion reaction.

C + O2 → CO2 Carbon dioxide is result of these reactions. Carbon dioxide produced is released to the air

through chimneys of power stations. Thus the cycle is completed and all carbon dioxide returns

to the atmosphere.

Green House Gases:

The sun sends energy to the earth in two forms, light and heat. Some of the heat energy reflects

back to the space, some however are trapped inside the Earth. This is caused by some gases and

it is called the greenhouse effect. The main greenhouse gases are carbon dioxide and methane.

• formation of carbon dioxide:

– as a product of complete combustion of carbon containing substances

– as a product of respiration

– as a product of the reaction between an acid and a carbonate

– from the thermal decomposition of a carbonate

Methane, the other greenhouse gas is formed by animals. When animals eat and digest their

food, methane gas is one of the waste products of this process. It is released to the atmosphere

by animals. When plants die and decompose over many years, methane gas is also produced.

The greenhouse effect poses a threat to the world now days. This is because greenhouse gases,

especially carbon dioxide, have increased in amounts in the atmosphere due to activity of

humans. Lots of fuel combustion is taking place around the world, increasing the levels of CO2,

while trees are being chopped off to made use of instead of leaving to replace CO2 with oxygen.

These activities cause an increase of the levels of CO2 in the atmosphere, which leads to more

heat trapping in earth. This rises the global temperature of the earth causing what’s called

global warming.

Global warming is the increase of the temperature of the earth due to the increase of levels of

greenhouse gases. Global warming has effects on the earth. To start with, it north and south

poles, which are made of ice, will start to melt raising sea levels. The sea temperature will also


rise causing death to marine lives. This is also accompanied by other natural disasters such as

hurricanes and heavy rains.

Humans could prevent this by reducing combustion of fossil fuels and leaving forests to live.

Rusting:

Rusting is the corrosion of iron as a result of reaction with oxygen from air and water. If iron

objects are left uncovered and exposed to air & water, iron will react with oxygen forming

hydrated iron oxide (also known as rust). Rust is a reddish brown flaky solid which will fall of

the object making it thinner and loses it its shape. Iron must come in contact with air and water

in order for rusting to happen. The formula of rust is Fe2O3. xH2O. Steel can also rust since it is

made up of mostly iron.

Rusting can become very dangerous in some cases. For example, bridges that cross rivers stand

on columns that are made of iron. The conditions of rusting are present in this case (Water from

the river and oxygen from the air). There is a risk that the columns will rust and collapse with

the whole bridge. In another case, ships are made of iron. Again, the conditions of rusting are

present (water from the sea and oxygen from the air). In fact, this situation is more critical

because sea water contains minerals that act as a catalyst to speed up the reaction of rusting.

There some available methods to prevent rusting. These methods are based on covering the iron

object with another substance to create a barrier between iron and oxygen and water so that

rusting does not take place:

Painting: The iron or steel object is painted all over. The paint creates the desired barrier to

prevent iron or steel coming in contact with air and water. This method is used in car bodies

and bridges.

Electroplating: The iron or steel object gets electroplated with another metal that doesn’t

corrode.

The object is usually electroplated with tin or chromium since they are very unreactive.

This method is used in food cans and car bumpers.

Sacrificial Protection: This method is based on the idea that metals that are higher than iron in

the reactivity series will react in preference to it and thus that metal is corroded and the iron is

protected. Metals usually used as protectors in this method are zinc and magnesium since they

are higher than iron in the reactivity series. In ships for example, zinc or magnesium bars are

attached to the iron base of the ship which is in contact with water and oxygen from air. But

rusting doesn’t take place since zinc or magnesium is the one that gets corroded. These bars

must be replaced from time to time because once they all get corroded, iron becomes

unprotected and rusts. This method is usually used in ships or bridge columns. The zinc or

magnesium bars do not have to completely cover the iron or steel because as long as they are

attached to each other the zinc or magnesium bars get corroded and not the iron.

Galvanisation: Galvanisation is a very reliable method for preventing rusting. It is basically

covering the whole object by a protective layer of zinc. This can be done either by

electroplating the object with zinc or dipping it into molten zinc. The zinc layer provides a

barrier that prevents iron or steel from coming in contact with air and water. The zinc gets

corroded instead iron thus protecting it. If the a part of the zinc coat falls off and the iron or

steel gets exposed to air and water, the bare part still doesn’t get corroded since it is protected

by sacrificial protection now.

Fertilisers

Chemicals applied to plants to improve their growth and increase the amounts of products such

as fruits, nuts, leaves, roots and flowers that they produce for us.

They work by supplying plants with the vital elements they need including


Nitrogen - in the form or nitrate (NO3- containing) salts;

phosphorous – in the form of phosphate (PO43-

containing) salts

and potassium (K+ containing) salts.

Displacement of ammonia from its salts

Ammonia (NH3) is a smelly gas.

One way to produce it is to react ammonium (NH4+) salts with an alkali (OH

-) eg:

NH4Cl + NaOH NH3 + H2O . + NaCl

The Haber process

The raw materials for the process of making ammonia are hydrogen and nitrogen.

Hydrogen is obtained by reacting natural gas (mostly methane) with steam, or

from cracking oil fractions.

Nitrogen is obtained from the air. Air is 78 per cent nitrogen and nearly all the rest is oxygen.

When hydrogen is burned in air, the oxygen combines with the hydrogen - leaving nitrogen

behind.

In the Haber process, nitrogen and hydrogen react together under these conditions:

a high temperature - about 450°C

a high pressure - about 200 atmospheres (200 times normal pressure)

an iron catalyst

In addition, any unreacted nitrogen and hydrogen are recycled. The reaction is reversible. In a

chemical equation, the symbol is used instead of an ordinary arrow if the reaction is

reversible:

This equation summarises the Haber process:

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Stage

1

Having obtained the hydrogen and nitrogen gases (from natural gas

and the air respectively), they are pumped into the compressor

through pipes.

Stage

2

The gases are pressurised to about 200 atmospheres of pressure

inside the compressor.

Stage

3

The pressurised gases are pumped into a tank containing beds of

iron catalyst at about 450°C. In these conditions, some of the

hydrogen and nitrogen will react to form ammonia.

Stage

4

The unreacted nitrogen and hydrogen, together with the ammonia,

pass into a cooling tank. The cooling tank liquefies the ammonia,

which can be removed into pressurised storage vessels.

Stage

5

The unreacted hydrogen and nitrogen gases are recycled by being

fed back through pipes to pass through the hot iron catalyst beds

again.

The reaction mixture contains some ammonia, plus a lot of unreacted nitrogen and hydrogen.

The mixture is cooled and compressed, causing the ammonia gas to condense into a liquid. The

liquefied ammonia is separated and removed. The unreacted nitrogen and hydrogen are then

recycled back into the reactor.


• To remember that carbon monoxide is poisonous (it binds to haemoglobin), think of the ‘nox’

in carbon monoxide as being short for noxious (poisonous). The effects of pollutant gases on

nature are often confused, as not all pollutant gases are acidic. Know the different effects of

carbon monoxide, sulphur dioxide and carbon dioxide

• A common error is to think that fume cupboards keep air away from a reaction. Fume

cupboards have a continuous airflow to allow poisonous vapours to escape through the fan

•You do not need to know all the details about water treatment. Filtration and chlorination are

the usual points examined.

•The separation of gases from the air is complex. You will only be asked questions about

boiling points and distillations – not about details of the distillation plant.

•It is a common error to suggest that sulfur rather than sulfur dioxide is responsible for acid

rain. Comments such as ‘sulfur dissolves in water to form acid rain’ are incorrect.

•The reactions in the catalytic converter are not well understood. The best equations to

remember are the reactions of nitrogen oxides with carbon monoxide to form nitrogen and

carbon dioxide.

•It is important that you do not muddle the effects of different pollutants: carbon dioxide and

methane are linked to global warming and sulfur dioxide is linked to acid rain.

•The two important regulating features of the carbon cycle are the uptake of carbon dioxide by

photosynthesis and the production of carbon dioxide during respiration.

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•If you are asked to choose two methods that prevent rusting, try not to choose two that are

similar. Don’t give answers such as ‘removing water and air’. These are explanations not

methods

•When you write equations for the formation of ammonium salts remember that no water is

formed as a product. For example: ammonia + sulfuric acid ammonium sulfate.

•You need to know the conditions used in the Haber process and why these particular

conditions are used by referring to the equilibrium reaction.


14 Sulfur

Sulphur is a non metal element in group 6 of the periodic table.

Sulphur has many useful properties which make it widely used in the industry.

Sources of Sulphur:

Sulphur is found in many places in the world in different forms. It usually exists in volcanic

regions in USA, Mexico and Sicily. Sulphur could also be obtained from some metal ores like

Copper pyrites (CuFeS2) and Blende (ZnS).

Properties of Sulphur:

In room temperature, sulphur is a yellow, brittle solid which doesn’t conduct electricity as it is a

non-metal. Sulphur is insoluble in water. It is able to react with both metals and non-metals.

Sulphur Dioxide:

Sulphur dioxide is the product of combustion of sulphur or sulphur-containing fuels. As you

have studied in the previous chapter, it is an air pollutant as it causes acid rain. However,

SO2 has important uses too:

Bleaching wood pulp for the manufacturing of paper

It is used as a food preservative as it kills bacteria

Manufacturing of Sulphuric acid

Contact Process (Manufacturing of Sulphuric Acid):

Sulphuric acid is one of the most important chemicals in the industry since it has a role in the

manufacturing of almost every product. Sulphuric acid is manufactured by a process called

Contact Process and it involves several steps:

1. Making the sulphur dioxide

2. Converting the sulphur dioxide into sulphur trioxide

3. Converting the sulphur trioxide into sulphuric acid

1. Making the sulphur dioxide

Sulphur is first burned in air producing sulphur dioxide:

S(s)+ O2(g)→ SO2(g)

2. Converting the sulphur dioxide into sulphur trioxide:

This is a reversible reaction, and the formation of the sulphur trioxide is exothermic.

2SO2(g) + O2(g) ⇌ 2SO3(g)

3. Converting the sulphur trioxide into sulphuric acid

This can't be done by simply adding water to the sulphur trioxide - the reaction is so

uncontrollable that it creates a fog of sulphuric acid. Instead, the sulphur trioxide is first

dissolved in concentrated sulphuric acid:


H2SO4(l) + SO3(g) → H2S2O7(l)

The product is known as fuming sulphuric acid or oleum.

This can then be reacted safely with water to produce concentrated sulphuric acid - twice as

much as you originally used to make the fuming sulphuric acid.

H2S2O7(l) + H2O(l) → 2 H2SO4(l)

The average percentage yield of this reaction is around 30%.

Properties & Uses of Sulphuric Acid:

Sulphuric acid is a very strong acid. It is a dibasic acid which means it every molecule of it

produces two hydrogen ions when it is dissolved in water. Sulphuric acid has some other unique

properties. For example, it is a dehydrating agent. This means it eliminates water from

compounds.

E.g.: CuSO4.5H2O CuSO4 + 5H2O

E.g.: C6H12O6 6C + 6H2O

It is also a drying agent. This means it removes water from mixtures. Don’t confuse that

dehydrating agent.


•Sulfuric acid has two hydrogen ions that can be replaced. Make sure that you remember this

when writing symbol equations for the reaction of sulfuric acid with metals, metal oxides and

metal carbonates.

•You will need to know the main reactions in the contact process and be able to write relevant

equations. You also need to know why the particular conditions of temperature and pressure are

used.


15 Carbonates

Carbonates are salts of carbonic acids (H2CO3). Carbonates are very useful salts, specially calcium carbonate (CaCO3).

Sources of Calcium Carbonate:

Calcium carbonate can be found in large amounts in the Peak District. It is found as a type of rocks called limestone near rivers.

Forms of Calcium Carbonate:

Limestone is not the only form of calcium carbonate. Marble and chalk are also other

forms of this valuable salt. Chalk is made of shells of marine algae. Marble on the

other hand, is a metaphoric rock made of limestone at high pressure.

Uses of Calcium Carbonate:

in the manufacture of iron and of cement

Manufacture of Lime:

One of the industrial uses of calcium carbonate is the manufacturing of lime from it. Lime is calcium oxide salt. This process takes place in a device called lime kiln and it is based on the thermal decomposition of calcium carbonate. Limestone is inserted in the kiln and heating starts. At the bottom of the kiln air is being blown in. this is also where lime is collected. The other product of this reaction, carbon dioxide gas, evolves and escapes at the top of the kiln.

CaCO3 ⇌ CaO + CO2

(Limestone) (Lime) (Carbon Dioxide)

Uses of Lime:

Lime can be used to neutralise soil acidity in farms. This is because it is a basic oxide. Slaked lime (Calcium hydroxide; Ca(OH)2) is also a basic oxide can be used as an alternative to lime for neutralising soil acidity. Another use of lime is neutralising sulphur dioxide waste in power stations. This is because sulphur dioxide is an acidic oxide while lime is a basic one. This process is called desulphurisation which you have studied earlier.


•It is a common error to suggest that oxygen is a reactant or product in the production of lime.

The reaction is a thermal decomposition. Oxygen does not react with calcium carbonate and

oxygen gas is NOT given off in the thermal decomposition.


16 Organic chemistry

16.1 Names of compounds

Alkanes share the same general formula:

Alkanes are saturated hydrocarbons. This means that their carbon atoms are joined to each

other by single bonds.

Alkenes are a homologous series of hydrocarbons that contain a carbon-carbon double bond.

The number of hydrogen atoms in an alkene is double the number of carbon atoms, so they

have the general formula .

Alcohols contain hydrogen and carbon but also possess one hydroxyl group (-OH). Their

general formula is CnH(2n+1)OH.

The names of alcohols end with ‘ol’, eg ethanol.

Carboxylic acids contain the carboxyl functional group (-COOH).

Carboxylic acids end in '-oic acid'.

Their general formula is CnH2O2.

16.2 Fuels

• Combustion

Fuels are substances that react with oxygen to release useful energy. Most of the energy is

released as heat, but light energy is also released.

About 21 per cent of the air is oxygen. When a fuel burns in plenty of air, it receives enough

oxygen for complete combustion.

Complete combustion

Complete combustion needs a plentiful supply of air so that the elements in the fuel react fully

with oxygen.

Fuels such as natural gas and petrol contain hydrocarbons. When hydrocarbons burn

completely:

the carbon oxidises to carbon dioxide

the hydrogen oxidises to water (remember that water, H2O, is an oxide of hydrogen)

Here are the equations for the complete combustion of propane, used in bottled gas:

propane + oxygen → carbon dioxide + water

C3H8 + 5O2 → 3CO2 + 4H2O

Incomplete combustion

Incomplete combustion occurs when the supply of air or oxygen is poor. Water is still

produced, but carbon monoxide and carbon are produced instead of carbon dioxide.

In general, for incomplete combustion:

hydrocarbon + oxygen → carbon monoxide + carbon + water

The carbon is released as soot.

Carbon monoxide is a poisonous gas, which is one reason why complete combustion is

preferred to incomplete combustion.

Gas fires and boilers must be serviced regularly to ensure they do not produce carbon

monoxide.

Here are the equations for the incomplete combustion of propane, where carbon is produced

rather than carbon monoxide:

propane + oxygen → carbon + water

C3H8 + 2O2 → 3C + 4H2O

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Fractional distillation of crude oil

Fractional distillation separates a mixture into a number of different parts, called fractions.

A tall fractionating column is fitted above the mixture, with several condensers coming off at

different heights. The column is hot at the bottom and cool at the top. Substances with

high boiling points condense at the bottom and substances with lower boiling points condense

on the way to the top.

Crude oil is a mixture of hydrocarbons. The crude oil is evaporated and its vapours condense

at different temperatures in the fractionating column. Each fraction contains hydrocarbon

molecules with a similar number of carbon atoms and a similar range of boiling points.

Oil fractions

The diagram below summarises the main fractions from crude oil and their uses, and the trends

in properties. Note that the gases leave at the top of the column, the liquids condense in the

middle and thesolids stay at the bottom.

The fractionating column

As you go up the fractionating column, the hydrocarbons have:

1. lower boiling points

2. lower viscosity (they flow more easily)

3. higher flammability (they ignite more easily).

• Name the uses of the fractions as:

– refinery gas for bottled gas for heating and cooking

– gasoline fraction for fuel (petrol) in cars

– naphtha fraction for making chemicals

– kerosene/paraffin fraction for jet fuel

– diesel oil/gas oil for fuel in diesel engines

– fuel oil fraction for fuel for ships and home heating systems

– lubricating fraction for lubricants, waxes and polishes

– bitumen for making roads

Other fossil fuels

Crude oil is not the only fossil fuel.

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Natural gas mainly consists of methane. It is used in domestic boilers, cookers and Bunsen

burners, as well as in some power stations.

Coal was formed from the remains of ancient forests. It can be burned in power stations. Coal

is mainly carbon but it may also contain sulfur compounds, which produce sulfur dioxide when

the coal is burned. This gas is a cause of acid rain. Also, as all fossil fuels contain carbon, the

burning of any fossil fuel will contribute to global warming due to the production of carbon

dioxide.

16.3 Homologous series

A homologous series is a family of compounds which have the same general formula and

have a similar molecular structure and similar chemical properties because they have the

same functional group of atoms (e.g. C=C alkene, C-OH alcohol or -COOH carboxylic acid).

Members of a homologous series have similar physical properties such as appearance,

melting/boiling points, solubility etc. BUT show trends in them e.g. steady increase in

melting/boiling point with increase in carbon number or molecular mass. The functional group is a group atoms common to all members of a homologous series that

confer a particular set of characteristic chemical reactions on each member molecule of the

series.

Characteristics of a homologous series:

-all the compounds fit the same general formula

-the chain length increases by 1 each time

-as the chain gets longer, the compounds show a gradual change in properties.

Structural isomers: have the same chemical formula, but different structures, they can be

straight or branched.

16.4 Alkanes

Alkanes are saturated hydrocarbons. This means that their carbon atoms are joined to each

other by single bonds.

This makes them relatively unreactive, apart from their reaction with oxygen in the air - which

we call burning or combustion.

Alkanes undergo a substitution reaction with halogens in the presence of light.

16.5 Alkenes

Bromine water is an orange solution of bromine. It becomes colourless when it is shaken with

an alkene. Alkenes can decolourise bromine water, but alkanes cannot.

The reaction between bromine and alkenes is an example of a type of reaction called

an addition reaction.

The bromine is decolourised because a colourless dibromo compound forms. For example:

ethene + bromine → dibromoethane

C2H4 + Br2 → C2H4Br2

Other addition reactions of alkenes:

Hydrogen can be added to a C=C double bond. This has the effect of ‘saturating’ the

molecule, and will turn an alkene into an alkane. For example: C2H4 + H2 → C2H6

If steam (H2O) is added to an alkene, an alcohol is made. For example: C2H4 + H2O →

C2H5OH

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Cracking

Fuels made from oil mixtures containing large hydrocarbon molecules are not efficient as they

do not flow easily and are difficult to ignite. Crude oil often contains too many large

hydrocarbon molecules and not enough small hydrocarbon molecules to meet demand. This is

where cracking comes in.

Cracking allows large hydrocarbon molecules to be broken down into smaller, more useful

hydrocarbon molecules. Fractions containing large hydrocarbon molecules are heated

to vaporise them. They are then either:

heated to 600-700°C

passed over a catalyst of silica or alumina

These processes break covalent bonds in the molecules, causing thermal

decomposition reactions. Cracking produces smaller alkanes and alkenes (hydrocarbons that

contain carbon-carbon double bonds). For example:

hexane → butane + ethene

C6H14 → C4H10 + C2H4

Some of the smaller hydrocarbons formed by cracking are used as fuels, and the alkenes are

used to make polymers in plastics manufacture. Sometimes, hydrogen is also produced during

cracking.

Alkenes can be used to make polymers.

Polymers are very large molecules made when many smaller molecules join together, end to

end. The smaller molecules are called monomers.

In general:

lots of monomer molecules → a polymer molecule

The polymers formed are called addition polymers.

Alkenes can act as monomers because they are unsaturated:

ethene can polymerise to form poly(ethene), also called polythene

propene can polymerise to form poly(propene), also called polypropylene

chloroethene can polymerise to form poly(chloroethene), also called PVC

Polymer molecules are very large compared with most other molecules, so the idea of a repeat

unit is used when drawing a displayed formula. When drawing one, you need to:

1. change the double bond in the monomer to a single bond in the repeat unit

2. add a bond to each end of the repeat unit

Addition polymerisation

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It can be tricky to draw the repeat unit of poly(propene). Propene is usually drawn like this:

It is easier to construct the repeat unit for poly(propene) if you redraw the monomer like this:

You can then see how to convert this into the repeat unit

16.6 Alcohols

The alcohols are a homologous series of organic compounds. They all contain the functional

group –OH, which is responsible for the properties of alcohols.

The first three alcohols in the homologous series are methanol, ethanol and propanol. They are

highly flammable, making them useful as fuels. They are also used as solvents in marker pens,

medicines, and cosmetics (such as deodorants and perfumes).

Ethanol is the alcohol found in alcoholic drinks such as wine and beer. Ethanol is mixed with

petrol for use as a fuel.

Ethanol from ethene and steam

Ethanol can be manufactured by the hydration of ethene. In this reaction, ethene (which comes

from cracking crude oil fractions) is heated with steam in the presence of a catalyst of

phosphoric acid (to speed up the reaction):

This reaction typically uses a temperature of around 300°C and a pressure of around 60–

70 atmospheres.

Notice that ethanol is the only product. The process is continuous – as long as ethene and steam

are fed into one end of the reaction vessel, ethanol will be produced. These features make it an

efficient process. However, ethene is made from crude oil, which is a non-renewable resource.

Ethanol can also be made by a process called fermentation.

Fermentation

During fermentation, sugar (glucose) from plant material is converted into ethanol and carbon

dioxide. This typically takes place at temperatures of around 30°C. The enzymes found in

single-celled fungi (yeast) are the natural catalysts that can make this process happen:

Unlike ethene, sugar from plant material is a renewable resource.

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Hydration of ethene v fermentation

16.7 Acids

Properties of carboxylic acids

Short carboxylic acids are liquids and are soluble in water. Longer carboxylic acids are solids

and are less soluble in water.

The boiling point of a carboxylic acid is higher than that of the alkane with the same number of

carbon atoms because the intermolecular forces are much stronger.

Carboxylic acids are weak acids, so they can donate a hydrogen ion(H+) in acid-

base reactions:

This means that they will react with carbonates to produce a salt, water and carbon dioxide:

They will also react with reactive metals to produce a salt and hydrogen.

Making a carboxylic acid

Ethanoic acid can be made by oxidising ethanol (which is an alcohol). In this case, oxidation

involves adding an oxygen atom and removing two hydrogen atoms. This can happen:

during fermentation if air is present

when ethanol is oxidised by an oxidising agent, such as acidified potassium

manganate(VII)

Making an ester

Esters occur naturally - often as fats and oils - but they can be made in the laboratory by

reacting an alcohol with an organic acid. A little sulfuric acid is needed as a catalyst.

The general word equation for the reaction is:

alcohol + organic acid → ester + water

For example:

methanol + butanoic acid → methyl butanoate + water

The diagram shows how this happens, and where the water comes from:

Fermentation Hydration of ethene

Type of raw

materials

Renewable (glucose from

plants)

Non-renewable (ethene from crude

oil)

Type of process Batch (stop-start) Continuous (runs all the time)

Labour A lot of workers needed Few workers needed

Rate of reaction Slow Fast

Conditions

needed

Warm (30°C), normal

pressure (1 atm)

High temperature (300°C) and high

pressure (60-70 atm)

Purity of product Impure (needs treatment) Pure (no by-products made)

Energy needed A little A lot

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So, to make ethyl ethanoate, you would need to react ethanol with ethanoic acid.

Different esters have different smells. These smells are often fruity.

Alcohol Organic acid Ester made Smell of ester

Pentanol Ethanoic acid Pentyl ethanoate Pears

Octanol Ethanoic acid Octyl ethanoate Bananas

Pentanol Butanoic acid Pentyl butanoate Strawberries

Methanol Butanoic acid Methyl butanoate Pineapples

16.8 Macromolecules

Macromolecules are large molecules built up from small units (monomers).

Different macromolecules have different units and/or different linkages

For example glucose (the small unit) can join together to make starch or cellulose (natural

macromolecules).

Examples of the small units:

-glucose

-amino acids

-fatty acids and glycerol

Examples of linkages:

-amide

-ester

Examples of macromolecules:

-starch

-protein

-lipids


16.8 (a) Synthetic polymers

Uses of polymers

Different polymers have different properties, so they have different uses. The table gives some

examples:

Polymer Typical use

Poly(ethene) Plastic bags ,bottles, gloves, cling film (low density), mugs,

bowls, chairs, dustbins (high density)

Poly(propene) Crates and ropes

Poly(chloroethene)

Water pipes, wellingtons, hoses and insulation on electricity

cables

Polystyrene Used as expanded polystyrene in fast-food cartons,

packaging, and insulation for roofs and walls

Teflon

Coated on frying pans to make them non-stick, fabric

protector, windscreen wipers, flooring

nylon ropes, fishing nets and lines, tents, curtains

Terylene Clothing (especially mixed with cotton), thread

Polymers have properties that depend on the chemicals they are made from and the conditions

in which they are made.

For example, there are two main types of poly(ethene) - LDPE, low-density poly(ethene),

and HDPE, high-density poly(ethene). LDPE is weaker than HDPE and becomes softer at

lower temperatures.

Modern polymers are very useful. For instance, they can be used as:

new packaging materials

waterproof coatings for fabrics (eg for outdoor clothing)

fillings for teeth

dressings for cuts

hydrogels (eg for soft contact lenses and disposable nappy liners)

smart materials (eg shape memory polymers for shrink-wrap packaging)

Pollution problems from plastics:

-choke birds, fish and other animals that try to eat them. Or they fill up the animals’ stomachs

so that they can’t eat

proper food, and starve to death.

-they clog up drains and sewers and cause flooding.

-they collect in rivers, and get in the way of fish. Some river beds now contain a thick layer of

plastic

-they blow into trees and onto beaches. So the place looks a mess. Tourists become put off.

• structure of the polymer product from a given alkene and vice versa



Condensation polymers

Some polymers are made via condensation polymerisation.

In condensation polymerisation, a small molecule is formed as a by-product each time a bond is

formed between two monomers. This small molecule is often water.

An example of a condensation polymer is nylon.

Condensation polymerisation involves linking lots of small monomer molecules together by

eliminating a small molecule. This is often water from two different monomers, a H from one

monomer, and an OH from the other, the 'spare bonds' then link up to form the polymer chain.

Nylon (a polyamide) is formed by condensation polymerisation, the structure of nylon

represented below where the rectangles represent the rest of the carbon chains in each unit.

For advanced molecular representations see Organic Nitrogen Compounds (A level Notes)

(3 units etc.)

This is the same linkage (-CO-NH-) that is found in linked amino acids in naturally occurring

macromolecules called proteins, where it is called the 'peptide' linkage.

Nylon-6,6

Terylene (a polyester) is formed by condensation polymerisation and the structure of

Terylene represented as

(3 units etc.)

This is the same kind of 'ester linkage' (-COOC-) found in fats which are combination of long

chain fatty carboxylic acids and glycerol (alcohol with 3 -OH groups, a 'triol').

Terylene (polyester) and nylon are good for making 'artificial' or

'man-made' fibres used in the clothing and rope industries.

In the manufacturing process the polymer chains are made to line up.

This greatly increases the intermolecular forces between the 'aligned'

polymer molecules and strong fibrestrands of the plastic can be made.

Although these are actually thermoplastic polymers, nylon and terylene can be drawn out

into thin strong fibres for use in clothing.

Some important structure, strength and 1D to 3D dimension concepts are in the Chemical

Bonding notes.

Nylon and polyester are typical synthetic fibres which have, in many cases, replaced cotton,

silk and wool fabrics in the clothing industry.


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They are cheap to make on an industrial scale compared to cotton from fields, silk from

silkworms and wool from sheep.

As well as being cheaper, the physical properties of synthetic fibres have several advantages

compared to their natural predecessors like cotton, silk and wool.

Compared to natural fibres, synthetic fibres tend to be ....

lighter - outdoor or indoor clothing,

more durable - harder tougher wearing fibres,

water-resistant - better water-proofed fabrics,

However, there are some disadvantages e.g.

they are not very breathable and sweat builds up making you feel uncomfortable.

A don't forget that silk fibres (for fabrics), rubber (for tyres and elastic objects) are very useful

natural polymers.

Wood, an extremely useful construction material, and is mainly a polymer mixture of cellulose

(a polymer of glucose) and lignin (with a rigid cross-linked structure).

The valuable crop of cotton (for fabrics) also has a molecular structure based on cellulose, in

fact its the purest form of cellulose that occurs naturally.

16.8 (b) Natural macromolecules

Food’s main constituents are proteins, fats and carbohydrates.

Proteins contain the same linkages (amide links) as nylon, but with different units.

Similarly, lipids and terylene both have ester links but different units.

The structure of a protein is:

In digestion proteins are broken down into amino acids (hydrolysis).

Fats are esters possessing the same linkage as Terylene (ester links) but with different units.

Soap is a product of the hydrolysis of fat. It is done using sodium hydroxide (as opposed to

acid, in digestion). The hydrolysis gives glycerol and the sodium salts of fatty acids. The salts

are used as soaps.

Complex carbohydrates: are a large number of joined sugar units (monosaccharide like

glucose).


Carbohydrates are a whole series naturally occurring molecules based on the elements carbon,

hydrogen and oxygen.

They are an important source of chemical energy in our diet.

e.g. the respiration reaction

glucose + oxygen ==> carbon dioxide + water

C6H12O6 + 6O2 ==> 6CO2 + 6H2O

Carbohydrates like glucose and fructose are used as sweeteners in food as well as sweets

themselves.

Historically the name 'carbohydrate' comes from the fact that all their formulae seemed to be

based on Cx(H2O)y BUT this is not the way to think of their formula.

They range from relatively small molecules called monosaccharide (means one basic unit),

or disaccharide (two basic units combined) to very large natural polymers or

macromolecules called polysaccharides (many units combined).

The formation of complex carbohydrates:

These are made of smaller carbon, hydrogen and oxygen based molecules combining together

e.g. the polysaccharides starch and cellulose are formed from glucose, molecular formula

C6H12O6.

Their formation can be described in terms of a large number of sugar units joined together

by condensation polymerisation

Note: Condensation polymerisation means the joining together of many small 'monomer'

molecules by eliminating an even smaller molecule between them to form the linkage.

e.g. HO-XXXXX-OH + HO-XXXXX-OH HO-XXXXX-O-XXXXX-OH + H2O etc.

n C6H12O6 ==> (C5H10O5)n + nH2O (where n is a very large number to form the natural

polymer)

The XXXXX or the [rectangles] below, represent the rest of the carbon chains in each unit

(more detail in the 3rd diagram below).

plus many H2O etc.


This diagram of starch or cellulose is in 'skeletal formula' style and both are polymers of

glucose - can you see the connection between each 'unit' and the structure of glucose itself?

The resulting natural polymer is called a polysaccharide.

Acid hydrolysis of complex carbohydrates (e.g.. starch) gives simple sugars.

This can be brought about by e.g. warming starch with hydrochloric acid solution to form

glucose.

(C5H10O5)n + nH2O ==> n C6H12O6 (where n is a very large number)

The hydrolysis products from polysaccharides can be analysed with paper chromatography as in

the case of amino acids.

We can digest long molecules like starch, though they have to be broken down by enzyme

action before the smaller molecules like glucose can be used in respiration.

However, we cannot digest cellulose because we don't have the enzymes to effect this process.

In digestion, the hydrolysis (Decomposition of a chemical compound by reaction with water,

such as the dissociation of a dissolved salt or the catalytic conversion of starch to glucose,

which can be accelerated by an acid or base) of starch happens in the mouth by the enzyme

amylase to make glucose.

In the lab, unless you have enzymes, you have to boil the complex carbohydrate (or proteins or

fats) in acid the products will be the following:

-starch → glucose

-proteins → amino acids

-fats → fatty acids and glycerol

But if hydrolysis is not complete, the macromolecules are not completely broken down.

So you get a mixture of molecules of different sizes for example for starch you get, glucose,

maltose (2 glucose units) and maltotriose (3 glucose units).

Chromatography can be used to identify the products and the substances. However, amino acids

and sugars are colourless when dissolved in water, so a locating agent is used. The substances

can be identified using the Rf values or by matching them with spots which are horizontal.

The fermentation of simple sugars to produce ethanol

Yeast is a microorganism containing an enzyme which will convert a sugar (glucose) solution

into carbon dioxide and alcohol (ethanol).

This process is called fermentation.

The word equation for fermentation is

glucose + yeast carbon dioxide + ethanol.

Carbon dioxide gas bubbles out of the solution into the air leaving a mixture of ethanol and water

Ethanol can be separated from the mixture by fractional distillation.

Fermentation must be carried out in the absence of air to make alcohol.

If air is present, ethanoic acid is made instead of alcohol.

Fermentation works best at a neutral or slightly acidic pH.

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http://www.gcsescience.com/w6-atmosphere-evolution-plant-oxygen.htm

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Fats and oils

Fats and oils are naturally-occurring esters. Fats are solid at room temperature, whereas oils are

liquids.

Vegetable oils

Vegetable oils are natural oils found in seeds, nuts and some fruit. The oil can be extracted. The

plant material is crushed and pressed and the oil, eg olive oil, is squeezed out.

Sometimes the oil is more difficult to extract and has to be dissolved in a solvent. Once the oil

is dissolved, the solvent is removed by distillation and impurities (such as water) are also

removed. This leaves pure vegetable oil, eg sunflower oil.

Structure of vegetable oils

Molecules of vegetable oils consist of glycerol and fatty acids.

The diagram shows how three long chains of carbon atoms are attached to a

glycerol molecule to make one molecule of vegetable oil.

The structure of a vegetable oil molecule

Fats and oils are esters possessing the same linkages as Terylene but with different monomer

units formed from long chain fatty acids and the 'triol'

alcohol glycerol , which has three C–O–H groups.

Glycerol is the alcohol plants and animals use to make oils and fats which are esters we use in

food and soaps.

Animals and plants combine glycerol and long chain fatty acids to make triglyceride esters –

fats from animals and oils from plants.

Most of them are esters of the tri–alcohol ('triol') glycerol (systematic name propane–1,2,3–

triol, but that can wait until AS–A2 level).

The carboxylic acids which combine with the glycerol are described as 'long–chain fatty acids'.

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The resulting ester is called a 'triester' or 'triglyceride' and they are the major components in

animal fat, vegetable oils, and processed fats like margarine etc..

'Traditional' soap is a product of the hydrolysis of fats from animals and vegetable oils from

plants

'Soapy' soaps (not modern detergents) are the sodium salts of long chain fatty acids formed by

heating fatty oils with concentrated alkalis like sodium hydroxide or potassium hydroxide to

hydrolyse them.

This is known as a saponification reaction and a typical equation is illustrated above and the

general word equation quoted below.

vegetable oil/animal fat + sodium hydroxide ==> soap molecule + glycerol

This reaction breaks the fat molecule down into one glycerol molecule (a triol alcohol) and

three sodium salts of the long chain carboxylic fatty acids that formed part of the original oil/fat

ester.


• Examiners are often very particular. One way to please them is to use the word ‘only’ in the

definition of a hydrocarbon i.e. the answer ‘a compound containing only carbon and

hydrogen’.

• Only one compound is formed in the reaction of ethene with steam. Remember, this is a

simple addition reaction (one compound formed from two or more substances) – a common

error is to say that hydrogen is also formed

• When trying to identify ‘cracking’ reactions from a set a reactions given, look out for one

molecule of reactant forming two or more molecules of product. Remember that cracking does

not involve oxygen

• ‘Clear’ does not mean ‘colourless’; when bromine is added to an alkene the colour change is

red-brown to colourless, not red-brown to clear

•When drawing the full structural formula of an organic compound you should show all atoms

and all bonds. Don’t forget that there is a bond in the alcohol functional group - O – H.

•When drawing alkenes make sure that there are not too many hydrogen atoms that form the

double bond. Check to see that each carbon has four bonds.

•Don’t get confused between petroleum and petrol. Petroleum is crude oil. Petrol, also known

as gasoline, is a fraction obtained when we distil petroleum.


•You do not have to remember the boiling range or typical number of carbon atoms in each

fraction and where they condense in the fractionating column.

•You will be expected to be able to balance symbol equations for the combustion of alkanes.

Remember to balance the oxygen.

•When describing cracking you must state that (i) large hydrocarbon molecules are broken

down to smaller ones (ii) and alkenes (iii) using a high temperature (iv) a catalyst.

•The test for an alkene is that it turns acqueous bromine colourless. Do not use the word ‘clear’

to. Aqueous bromine stays the same yellow or orange colour when an alkene is added. A comm

mean colourlesson error is to write ‘no observations’.

•When writing a symbol equation for the combustion of an alcohol, when you balance the

oxygen, remember that the alcohol contains oxygen too.

•Remember that in naming the carboxylic acids the carbon atom of the –COOH group is

included. So CH3COOH is ethanoic acid because compounds with two carbon atoms have

names beginning with ‘-eth’.

•When plastics are burned, poisonous or toxic gases are given off– not just harmful gases. A

common error is to suggest that sulfur dioxide is given off when plastics burn. Few plastics

contain sulfur.

•When writing the formula for an addition polymer don’t forget: (i) the double bond changes to

a single bond and (ii) to include the continuation bonds.

•A common error in writing formulae for polyamides and polyesters is to write all the bonding

atoms in the same direction when the monomers each have only one type of functional group, is

wrong

•You should be able to recognize the repeating units in proteins as NHCOCH(R) and that this

repeats along the chain.

•The full structure for a fat and the equation for soap making do not need to be remembered.

Sufficient information will be given to help you answer questions.

•You do not need to know the structure of carbonhydrates but it is important to know how

hydrolysis breaks down complex carbonhydrates using simplified formulae.

•It is a common error to suggest that oxygen is required for the fermentation of glucose to

ethanol.


GLOSSARY

acid = any substance that produces hydrogen ions, H+, when dissolved in water

acidic solution = a solution with a pH less than 7

acid rain = rain with a pH less than 5.6; acid rain has been made more acidic than normal

rain because sulfur oxides and nitrogen oxides have dissolved in it; acid rain causes damage to

buildings made from limestone, damages metal structures, kills fish, damages leaves in trees so

they photosynthesise less

acidic soil = soil with a pH less than 7

activation energy = minimum amount of energy needed to start the reaction/for a successful

collision.

actual yield = the amount of product obtained when carrying out a reaction

addition reaction = a reaction in which atoms are added to an unsaturated carbon compound;

the atoms are added using the double bond as one of the double bonds breaks and is used to

make two new bonds, e.g. alkenes and halogens

addition polymer = polymer formed by addition polymerization; adding many unsaturated

monomers using double bonds

addition polymerization = the joining together of many unsaturated monomer molecules

(double bonds) to form a long molecule; new monomers are added to the chain at the double

bonds

alcohol = a homologous series of organic compounds which has -OH as its functional group;

ethanol is a member of this homologous series

alkanes = a homologous series of hydrocarbons which are saturated as they have only single

bonds between the carbon atoms

alkenes = a homologous series of hydrocarbons which are unsaturated as they have at least 1

double bond somewhere in the chain

allotrope = different forms of the same element e.g. diamond, graphite and the fullerenes are

allotropes of carbon

alloys = mixture of a metal and small amounts of other metals and non-metals, made to have

certain improved properties eg harder, stronger, increased resistance to corrosion, increased

heat or electrical resistance

alkali = any base which is soluble in water

alkali metal = any metal in group 1 of the Periodic Table, most reactive metals

alkali solution = a solution with a pH larger than 7

anions = negative ions; attracted to anode

anode = positive electrode in electrolysis

arrangement = how particles are positioned compared to each other e.g. close together, far

apart, in fixed positions

atom = the smallest particle that can exits of an element

atomic number = number of protons in the nucleus of an atom, determines the order and

place of each element in the Periodic Table

avogadro’s constant = 6.02 x 1023

balanced equation = numbers of atoms are the same on either side of the equation (any

equation should be balanced as in any chemical reaction particles are only re-arranged and are

not destroyed or created); also shows the ratio in which reactants react and products are

produced during a chemical reaction

base = a substance which can neutralise an acid to make a salt and water examples: metal

oxides, metal hydroxides,

bauxite = ore containing aluminium oxide from which aluminium is extracted

blast furnace = a furnace used for getting iron from iron oxide with the help of carbon

boiling = a process during which a liquid changes into a gas as its particles gain more energy

and move a lot faster and also much farther apart from each other. further from gas to liquid;

only happens at the boiling temperature as opposed to evaporation

brine = concentrated sodium chloride solution

catalyst = a substance which speeds up a reaction but which remains unchanged at the end of

the reaction

catalytic converter = a piece of equipment which is part of the exhaust of a car and which

changes nitrogen oxides into nitrogen before they are released into the atmosphere


cathode = negative electrode in electrolysis

cation = positive ion ; attracted to cathode

chain length = number of carbon atoms one after the other in an organic compound

chemical bond = electrostatic attraction between atoms or ions

chemical property = how it reacts

chromatogram = the result of a chromatography

chromatography = a separating technique which uses the difference in solubility in a given

solvent between the different parts of a mixture to separate them;

combustion = burning, the reacting of a substance with oxygen, exothermic

complete combustion = combustion in sufficient oxygen which in the case of hydrocarbons

produces carbon dioxide and water

compound = a pure substance made from two or more different atoms joined together

chemically

concentration = the number of moles of per liter of solution; tells us how much solute is

dissolved inn a solvent

condensation = a process during which a gas changes into a liquid because its particles are

having less energy, slow down and come much closer together

condensation polymer = a long molecule formed by condensation polymerization e.g. nylon

condensation polymerization = the joining together of many of two different monomer

molecules to form one single long molecule during which a small molecule is removed for each

link between the monomers.

covalent bond = force of attraction between a pair of shared electrons and the nucleii of both

atoms

cracking = the breaking down of long-chain alkanes into smaller alkanes and alkenes using a

catalyst and heat (500 C)

crude oil (or petroleum) = a mixture of organic compounds formed, as a result of high

temperatures and pressures, from the remains of living plants and animals which died millions

of years ago; a fossil fuel

crystallisation = the forming of crystals from a saturated solution

decomposition = breaking down a compound into simpler substances

delocalised electrons = electrons that can move between atoms; they are not part of 1 atom

diamine = a type of amine with exactly two amino groups

diatomic = 2 atoms only

dicarboxylic acid = organic compounds that contain two carboxylic acid functional groups.

diffusion = the movement of particles by which different substances mix as a result of the

random motion of each of their particles

displacement reaction = a reaction in which a more reactive metal or halogen takes the place

of a less reactive metal or halogen in its compound

distillate = the liquid obtained from distillation; the liquid which has evaporated and

condensed

distillation = a separating technique in which a mixture is heated, the substance with the

lowest boiling point evaporates and is condensed back to liquid form

ductile = can easily be drawn into wires, what metals are

endothermic = absorbs energy

electrical conductivity = conducts electricity for which it needs mobile charged particles

electrodes = rods of ususally carbon which are used to make elctrical contact with the

electrolyte

electrolysis = a reaction which uses electricity to decompose a compound

electrolyte = an ionic compound or acid which conducts electricity (molten or in solution)

and which is decomposed as it conducts

electrolytic cell = a beaker with an electrolyte, 2 electrodes, a power supply and leads which

changes electrical energy into chemical energy

electron = a sub-atomic particle which has a negative charge and no relative mass

electronic configuration = the number of electrons on each energy level in an atom

element = a pure substance that consists of 1 type of atom only

empirical formula = the formula which gives the most simple ratio of atoms/ions in a

molecule/formula unit

http://en.wikipedia.org/wiki/Polyamine

http://en.wikipedia.org/wiki/Amino_group

http://en.wikipedia.org/wiki/Organic_compound

http://en.wikipedia.org/wiki/Carboxylic_acid

http://en.wikipedia.org/wiki/Functional_group


equilibrium = is reached when the forward reaction and reverse reaction are going on at the

same time; at this point the amount of reactant or product does not change.

evaporation = a process during which a liquid changes into a gas as some of its particles at

the surface gain more energy, move a lot faster and farther apart from each other and eventually

escape from the liquid; happens at any temperature between melting and boiling point.

exothermic = releases/gives out energy

fermentation = the changing of sugars dissolved in water into alcohol and carbon dioxide by

the enzymes in yeast at a temperature of between 30 C to 40 C.

filtrate = the liquid/solution that goes through the filter paper

fixed positions = particles in a solid cannot move from their positions because of the strong

forces of attraction

forces of attraction = forces which hold/pull particles together

forward reaction = the reaction which produces the products

fraction = a group of substances which has a specific boiling point/range/condenses at similar

temperature (because they have a similar number of carbon atoms in them);

fractional distillation = crude oil is heated to evaporate most components which then

condense back at different levels in the fractionating column because they have differing

boiling points;

freezing = process during which a liquid changes into a solid as its particles lose energy, slow

down and come closer together again

fuel = a substance that can release a lot of energy e.g. by burning

gas = a state of matter in which particles are far apart, have a lot of energy and move fast and

randomly

galvanising = the coating of steel or iron by zinc to protect it from rusting

giant structure = a structure in which a very large number of atoms or ions are joined

together strongly and continuously in all 3 directions; a large network of particles

group = vertical column in Periodic Table

half equation = equation showing what goes on at each electrode in electrolysis

halogen = any element from group 7 in the Periodic table

homologous series = a group of organic compounds which all have the same general

formula, similar chemical properties because they have the same functional group, have a

gradual trend in physical properties, and differ by one CH2 unit.

hydrocarbon = a compound which has carbon and hydrogen only

incomplete combustion = burning in not enough oxygen

indicator = any chemical which can change colour when added to different chemicals,

usually acids and bases

inert = very unreactive

inert gases = gases in group 0

intermolecular forces = weak forces of attraction between molecules

ion = a charged atom or group of atoms (which has become charged because it has either lost

or gained an electron(s))

ionic bond = strong electrostatic attraction between two oppositely charged ions, formed

between metals and non-metals

isomers = compounds with the same molecular formula but different structures or displayed

formula and therefore different properties

isotopes = atoms with the same number of protons and electrons but different number of

neutrons; same mass number but different mass number

lattice = regular 3-dimensional arrangement of the particles (atoms, ions or molecules)

limestone = calcium carbonate

liquid = a state of matter in which particles are close together but in a disorderly arrangement,

they can move past one another and have energy to move from their positions

lubricant = an oily but soft substance used to reduce friction between two moving surfaces

malleable = easily shaped without breaking, what metals are

mass number = the total number of protons and neutrons in the nucleus of an atom

melting = a process during which a solid changes into a liquid as its particles have gained

more energy and move from their positions and past one another into an irregular arrangment

metallic bond = attraction between positive metal ions and delocalised (mobile ‘sea’ of

electrons electrons


metallic character = behaves like a metal, gives away electron (s) when it reacts to form a

positive ion, conducts, shiny, malleable

mixture = 2 or more substances mixed together which have not reacted and which are

therefore easily separated by physical processes like evaporation/distillation/filtration

molar mass = the actual mass of 1 mole or 6.02 x 1023 particles (atoms, ions, molecules or

formula units) of that substance

molar volume = the volume of 1 mole of a gas = 24L at rtp

mole = the name given to a certain number and that number is 6.02 x 1023.

molecular formula = shows the type of atoms/ions and their number/molar ratio in a

molecule/formula unit

molecule = a particle made up of 2 or more atoms held together by covalent bonds

monomer = a small molecule which can be joined together to make a long molecule called a

polymer; a monomer must have a double bond or a functional group at either end

movement = how particles move e.g. fast, vibrate

neutron = a sub-atomic particle with no charge and which is in the nucleus and has a relative

mass of 1

neutralisation = a reaction between an acid and a base to produce water and a salt and

sometimes also carbon dioxide

noble gas = any element form the last group in the Periodic Table

noble gas electronic configuration = the way in which electrons are arranged in the noble

gas atoms which is that they have their outer shell full! This 2 electrons in the helium atom and

8 in the other noble gases

ore = a mixture of rock which contains a useful chemical

organic compounds = compounds that have the element CARBON in it

oxidation = a reaction during which a substance gains oxygen; oxygen is added to the

element or compound increasing its mass; also a reaction during which a substance loses an

electron

oxide = a compound which ends with oxygen

oxidizing agent = a chemical which oxidises another chemical; it loses oxygen/gains

electrons and becomes reduced

oxide layer = layer of an oxygen compound

% yield = the actual yield expressed as percentage of the theoretical yield

period = horizontal row in the Periodic Table

periodic trends = gradual changes in properties of the elements in the same period

petroleum = a mixture of organic compounds formed, as a result of high temperatures and

pressures, from the remains of living plants and animals which died millions of years ago;

contains fossil fuels.

pH = a number between 1 and 14 which tells us how strong or how weak an acid or alkali is

pH scale = a scale running from 1- 14 used to show how acid or alkaline a substance is

physical property = properties like melting and boiling point, volatility, conductivity,

appearance, colour

poly(ethene) = polymer made from polymerising ethene molecules - addition polymer

polymer = a large molecule made from many small molecules that have been joined together;

each polymer is made up of many repeated units

polymerisation = a chemical reaction in which many small molecules or monomers are

joined together to form a long molecule called a polymer

position of equilibrium = gives an idea of how much reactant or product there is at

equilibrium

precipitation = a reaction between 2 salt solutions which produces an insoluble salt which

sinks to the bottom of the test tube

precipitate = insoluble solid formed during a reaction

product = substance on right hand of equation

pure substance = a single chemical element or compound which melts and boils at fixed

temperatures

rate of a reaction = amount of change in a reactant or product over a period of time; tells us

how fast a reaction is going

reactant = substance on left side of equation

reactivity = refers to the ease with which a substance reacts with other substances


reactivity series = a list of metals with the most reactive metal first based on results from

experiments

redox = a reaction during which both a chemical is oxidised and another is reduced

reducing agent = a chemcial which reduces another chemical; it gains oxygen/loses electrons

and becomes oxidised

reduction = a reaction during which a susbstance loses oxygen and has its mass decreases;

also a reaction during which a substance gains electrons

relative atomic mass = the mass of an atom as compared to 1/12th of the mass of a carbon-

12 atom; it is also the average mass of all isotopes

relative molecular mass = the sum of the relative atomic masses (multiplied by the number

of times they are in the molecule) of the atoms in the molecule

relative formula mass = the sum of the relative atomic masses (multiplied by the number of

times they are in the formula) of the atoms or ions in the giant structure

residue = the insoluble part that remains behind in the filter paper during a filtration or what

is left in the flask

reverse reaction = reaction which changes products back into reactants

reversible reaction = a reaction during which products are made but are also changed back

again into reactants

rust = a loose orange brownflaky layer of hydrated iron oxide

sacrificial protection = method of rust protection in which blocks of more reactive metal are

attached to irn; the more reactive metal react with the air and water instead of the iron

saturated solution = a solution which contains as much solute as possible

saturated organic compound = each carbon atom in the organic compound has made 4

single covalent bonds

simple molecular substance = substance made up of individual molecules held together by

covalent bonds and has weak intermolecular forces between these molecules

solid = a state of matter in which particles are close together and in a regular arrangement,

can only vibrate in fixed positions and have little energy

solvent = a liquid that does the dissolving

solvent front = the height the solvent goes up to on the chromatography paper

solute = a solid which dissolves

solution = a mixture made by dissolving a solute in a solvent

steel = an alloy of iron with other elements

sub-atomic particle = very small particles from which atoms are made: electrons, protons

and neutrons

sublimation = a process during which a solid changes directly into a gas because its particles

have alot more energy, move around very fast and are very far apart.

system = the reactants and products of a reaction

theoretical yield = the amount of product you should obtain according to the balanced

equation and calculations

thermal decomposition = breaking down of a compound by hetaing it

transition element = metal in the transition block of the Periodic Table

universal indicator = a mixture of indicators used to measure the pH because it goes

different colours

unsaturated organic compound = has at least one double bond; decolourizes brown bromine

water

valency = the combining power of an atom or group of atoms; in an ionic compound the

valency of an ion is its charge; in a molecule the valency of an atom is the number of bonds it

makes

valency electrons = the electrons on the most outer shell;

vapourise = change from liquid into gas

vibrate = move forwards and backwards but in the same fixed position

volatile = vapourises easily, low boiling point

word equation = an equation in which the names of the chemicals are used


EXAMINER TIPS

PAPER 1

• Some questions may ask you to choose a combination of things in order to select the correct

answer – understand exactly what is required before you start to answer

• Within a single question, use a pencil to cross out the choices that are clearly incorrect, then

choose between the others

GENERAL TIPS FOR PAPERS 2, 3, 5 AND 6

• Read the question correctly. For example, if the question says ‘give two observations apart

from temperature change’, don’t include temperature change in your answer

• Check for contradictions within your answer. For example, a common error is to write ‘a

white insoluble precipitate dissolves’ (6.5(a)(i)) Show any workings

• In any calculation, the final answer should be to the correct number of significant figures –

generally the same as the data. You may be penalized if you write an excess number of

significant figures e.g. 1.257487 instead of 1.26

• Know your syllabus statements and definitions exactly – use the Revision Checklist on the

website. Don’t add your own ideas to the statements. For example, the syllabus statement on

batteries says ‘they are portable’, meaning they can be easily carried around: an answer such as

‘they are small’ may not be accepted, as something can be small yet heavy

• If asked to ‘describe what you would observe’, write down what you see, hear or feel (e.g.

‘the test tube gets hot’). A common mistake is to write something like ‘a gas is given off’ or

‘copper is deposited’; these are not observations, these are conclusions

• If asked to ‘describe what you would see’, don’t note observations about sounds or

temperature

• Learn your definitions! Questions such as “what is a compound?” or “Define the mole.” are

often poorly answered. Define does not mean ‘give an example of.’

• When drawing diagrams:

(i) make sure they fill the space given on the paper and are LABELLED

(ii) when drawing apparatus for gas measurement, make sure that the gas cannot

escape. For example, don’t draw a gas syringe with the plunger much smaller than

the syringe barrel – this is a common error

• When asked to give examples, give the number requested by the examiner. For example, if

asked to give two examples, do not give three – if one is incorrect, you may lose a mark. If a

question asks for a single use for a substance don’t write a list – the examiner will think you are

‘playing safe’ and you won’t get the mark

• If you have to tick boxes to answer a question, make sure that you tick the correct number –

don’t assume that a single answer is always required

• In chemistry, when plotting a graph of reaction rate, you must draw a curve of best fit through

your prints. Lines drawn with a ruler from point to point will not get a mark

• Look out for ‘hidden words’ in questions such as ‘which of the following is a gas containing

diatomic molecules?’ Many students focus on one or two words, and might forget ‘gas’.

Underline key words and read the question slowly

• Avoid vague statements. For example, if the question asks about the use of graphite, the

answer ‘graphite is used for electrodes in electrolysis’ is appropriately specific. ‘Graphite is

used in electrolysis,’ is too vague


TIPS ON PRACTICAL PAPERS

• Makes sure that you know the accuracy to which you can read burettes, measuring cylinders,

etc.

• If asked to describe the appearance of a substance, remember that there are generally two

points to be made, the state and the colour

• When making practical observations, use the words ‘precipitate’ rather than ‘cloudy’ if you

cannot see through the test tube on adding one aqueous solution to another. Don’t forget the

colour

• ‘Test the pH’ means ‘give the pH number’, not just say whether something is acidic or

alkaline

•In (a) make sure that full and correct names are used. ‘Cylinder’, ‘stand’ and ‘spoon’ are not

precise. ‘Measuring cylinder’, ‘tripod’ and ‘spatula’ are the correct names of apparatus.

In (b) use a ruler and a sharp pencil and label the diagram clearly. A filter funnel and filter

paper both need to be included in the answer.

•Common mistakes are to label the box with the wrong acid. The solid is often incorrectly

labeled as sodium sulfate, rather than sodium sulfite. In the answer to (b) an arrow needs to be

positioned with its point touching the flask underneath the solid. Vague answers in (c) will not

score credit, e.g. ‘There is no lid on the collecting vessel.’ Identifying clearly that the gas

should be collected through water are the correct conclusions drawn from the supplied

information.

•Note the readings after checking the scale used in the diagrams.

•A common error would be to misread the temperatures recorded when 10cm3 and 40cm

3 of Y

have been added. Incorrect readings would be ‘36C’ and ‘46C’.

•This question involves applying experience of common practical procedures to an unfamiliar

situation. It also tests knowledge and understanding of chromatography. A diagram in (d)

showing the solvent at the correct level and three separate colourings would indicate the ability

to present information in a clear, logical form. When drawing chromatography apparatus, make

sure that the place where you put the original spot of colour is above the level of the solvent.

•The answer in (b) show that the knowledge of cation tests is good. The student correctly

describes the effect of aqueous ammonia and sodium hydroxide on a solution containing Zn2+

(aq).

•In the answers to (b) (i) and (ii) the examiners would prefer you to use the word ‘colourless’

instead of ‘clear’ when referring to a solution that is not coloured.

•The anion tests in (e) and (f) are also recognized. However, the use of cobalt chloride to test

for the presence of water was confused with the test for chlorine gas.

•The standard of answers varied widely from excellent to very poor. Poor answers involved :

mixing the fuel with the water and measuring the temperature rise

heating the fuel directly and measuring the temperature rise

failure to ensure a fair test

reference to the diagram with no detailed method.

Failure to show how a comparison of the results would indicate which fuel produces more

energy was also common.


•The response shows the ability to suggest suitable techniques and apparatus for the

investigation. All measurements and observations to be made are clearly recorded. The idea of a

fair test is clearly realized and the comparison of the results to draw appropriate conclusions

made. This student also notes safety precautions and suggests repeating the experiment to check

reliability.

Igcse2

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