Ap chem unit 7

ATOMIC STRUCTURE AND PERIODICITYAP Chem: Unit 7

ELECTROMAGNETIC RADIATION


One of the ways that energy travels through space is by electromagnetic radiation.

light from the sun X-rays microwaves

WAVE CHARACTERISTICS Wavelength (λ) – is the distance

between two consecutive peaks or troughs in a wave.

Frequency (ν) – is the number of waves (cycles) per second that pass a given point in space. units – hertz or waves/sec (s-1)

Speed (c) – all types of electromagnetic radiation travel at the speed of light.2.9979 x 108 m/sc = λν

THE NATURE OF MATTER

WAVE AND PARTICLE DUALITY

Planck found that matter could only absorb or emit energy in whole number multiples of the quantity hν.

h is Planck’s constant = 6.626 x10-34 Js

ΔE = hν Transfer of energy is not continuous

but is quantized and can occur only in discrete amounts called quantum. Thus energy has particle properties as well as wave properties.

EINSTEIN

WAVE AND PARTICLE DUALITY

Einstein proposed that electromagnetic radiation was also quantized and could be viewed as a stream of “particles” called photons.

Ephoton = hv = hc/λ

THE PHOTOELECTRIC EFFECT

The photoelectric effect refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it.

1. No electrons are emitted by a metal below a specific threshold frequency (vo)

2. For light with frequency lower than the threshold frequency, no electrons are emitted regardless of intensity of the light.


3. For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light.

4. For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases directly with frequency of the light.


These observations can be explained by assuming that electromagnetic radiation is quantized (consists of photons), and that the threshold frequency represents the minimum energy required to remove the electron from the metal’s surface.

Minimum energy required to remove an electron = Eo = hvo

KEelectron = ½ mv2 = hv – hvo

PLANCK AND EINSTEIN CONCLUSIONSEnergy is quantized. It can occur

only in discrete units called quanta.

Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. This phenomenon is sometimes referred to as the dual nature of light.

WAVE PARTICLE DUALITY

The main significance of the equation E = mc2 is that energy has mass.

m = E/c2

LOUIS DE BROGLIE (1892-1987)

Since light which previously was thought to be purely wavelike, was found to have certain characteristics of particulate matter. But is the opposite also true? Does matter have that is normally assumed to be particulate exhibit wave properties?


de Broglie’s equation allows us to calculate the wavelength for a particle:

λ =h

mv

DE BROGLIE’S PROOF


Conclusion: Energy is really a form of matter, and all matter shows the same types of properties. All matter exhibits both particulate and wave properties.

THE ATOMIC SPECTRUM OF HYDROGEN

SPECTRUMA continuous spectrum results

when white light passes through a prism and all wavelengths (colors) are shown.

An emission spectrum produces only a few lines of color that is limited to discrete wavelengths produced by an atom. This is called a line spectrum and is specific to each atom.

HYDROGEN LINE SPECTRUMThe significance of the line

spectrum is that it indicates that only certain energies are allowed for the electron in the hydrogen atom. In other words the energy of the electron in the hydrogen atom is quantized

HYDROGEN LINE SPECTRUM

THE BOHR MODEL

NIELS BOHRBohr developed a

quantum model for the hydrogen atom that allowed for only specific energy levels around the atom that corresponded with specific radii.

NIELS BOHR (1885-1962)The most important equation to

come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.

Z is the nuclear charge, n is the energy level.

E =−2.178x10−18 JZ2

n2

⎛

⎝⎜⎞

⎠⎟

NIELS BOHR (1885-1962)The most important equation to come

from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.

the negative sign calculates a lower energy closer to the atom, not the radiation of negative energy.

E =−2.178x10−18 JZ2

n2

⎛

⎝⎜⎞

⎠⎟

EXAMPLE

What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?

ni=6; nf=1; Z=1 (hydrogen nucleus contains a single proton)

EXAMPLE

What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?

E6 =−2.178x10−18 J12

62

⎛

⎝⎜⎞

⎠⎟=−6.050x10−20 J

E1 =−2.178x10−18 J12

12

⎛

⎝⎜⎞

⎠⎟=−2.178x10−18 J

EXAMPLE

ΔE=Ef – Ei = E1 – E6=-2.117 x 10-18J

The negative sign for the change in energy indicates that the atom has lost energy and is now more stable. This loss of energy produces a photon.

EXAMPLE

What is the corresponding wavelength for the energy produced from the electron jump?

E = -2.117 x 10-18J

E =hv=hcλ

⎛⎝⎜

⎞⎠⎟

9.383x10-8 m

BOHR MODEL CONCLUSIONS The model correctly fits the quantized

energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electrons.

As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state. As the electron is brought closer to the nucleus, energy is released from the system.

BOHR MODEL CONCLUSIONS

ΔE = E final − Einitial =

−2.178x10−18 J1

n final2−

1

ninitial2

⎛

⎝⎜

⎞

⎠⎟

BOHR MODEL CONCLUSIONSThe energy levels calculated by

Bohr closely agreed with the values obtained from the hydrogen emission spectrum but does not apply well to other atoms. The Bohr’s model is fundamentally incorrect but is very important historically because it paved the way for our current theory of atomic structure.

THE QUANTUM MECHANICAL MODEL OF THE ATOM

QUANTUM MECHANICS

Quantum Mechanics or Wave Mechanics were developed by three physicists: Heisenberg, de Broglie, and Schrodinger.

Emphasis was given to the wave properties of the electron.

The electron bound to the nucleus behaves similar to a standing wave.

QUANTUM MECHANICS Like a standing

wave, electrons can travel in patterns that allow for a common node. In other words, wave patterns around the nucleus must be in whole number wave patterns. But their exact movement is not known.

HEISENBERG UNCERTAINTY PRINCIPLEThere is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. This limitation is small for large particles but substantial for electrons.

Δx • Δ mv( ) ≥h

4π

PROBABILITY DISTRIBUTION

A probability distribution is used to indicate the probability of finding an electron in a specific position.

Electron density map

Radial probability distribution

PROBABILITY DISTRIBUTION

For the hydrogen 1s orbital, the maximum radial probability occurs at a distance of 5.29x10-2nm or .529Å from the nucleus. This is the exact radius of the innermost orbit calculated in the Bohr Model.

The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability

QUANTUM NUMBERS

QUANTUM NUMBERS

Each orbital is characterized by a series of numbers called quantum numbers, which describe various properties of an orbital:

1. Principal quantum number (n)- has integral values : 1,2,3,4. It describes the size and energy of the orbital. Energy Level

QUANTUM NUMBERS

2. Angular momentum quantum number (l) – has integral values from 0 to n-1. This is related to shape of the atomic orbitals. Sublevelo l =0 is so l =1 is po l =2 is do l =3 is fo l =4 is g

QUANTUM NUMBERS

3. Magnetic quantum number (ml) has values between l and – l , including 0. The value of ml is related to the orientation of the orbital in space. Axis designation

QUANTUM NUMBERS

4. Electron spin quantum number (ms)- can only have one of two values, +½, -½. Electrons can spin in one of two opposite directions.

QUANTUM NUMBERS

In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml , ms). This is called the Pauli exclusion principle; an orbital can only hold two electrons, and they must have opposite spins.

QUANTUM NUMBERS

ORBITAL SHAPES AND ENERGIES

S ORBITALSThe s orbitals have

a characteristic spherical shape and contain areas of high probability separated by areas of zero probability. These areas are called nodal surfaces, or nodes.

S ORBITALSThe number of

nodes increases as n increases. The number of nodes = n - 1.

P ORBITALS P orbitals each have two lobes separated

by a node at the nucleus. The p orbitals are labeled according to the axis of the xyz coordinate system along which the lobes lie.

P ORBITALSCross section of electron probability of a p

orbital

D ORBITALS The five d orbitals first occur in energy

level 3. They have two fundamental shapes. Four of the orbitals (dxz, dyz, dxy, and dx2-y2) have four lobes centered in the plane indicated in the orbital label. dx2-y2 lie along the x and y axes and dxy lie between the axes. The fifth orbital dz2 has a unique shape with two lobes along the z axis and a belt centered in the xy plane.

D ORBITALS

F ORBITALSThe f orbitals first occur in level 4 and

have shapes more complex than those of the d orbitals. These orbitals are not involved in the bonding in any of the compounds that we will consider.

F ORBITALS

ORBITAL ENERGIES For the hydrogen atom, the

energy of a particular orbital is determined by its value of n. Thus all orbitals with the same value of n have the same energy – they are said to be degenerate.

POLYELECTRONIC ATOMS

POLYELECTRONIC ATOMSPolyelectronic atoms are atoms with

more than one electron. To look at these atoms, three energy contributions must be considered:

Kinetic energy of the electrons as they move around the nucleus.

The potential energy of attraction between the nucleus and the electrons.

The potential energy of repulsion between the two electrons.


Since electron pathways are unknown, dealing with the repulsions between electrons cannot be calculated exactly.

This is called the electron correlation problem.


The electron correlation problem occurs with all polyelectronic atoms. To deal with this, we assume each electron is moving in a field of charge that is the net result of the nuclear attraction and the average repulsions of all the other electrons.

In other words,…..


A valence electron is attracted to the highly charged nucleus and still repelled by the other ‘inner’ electrons. The net effect is that the electron is not bound nearly as tightly to the nucleus as it would be if it were alone.

This is a screened or shielded affect.


Because of this shielded affect. orbitals within a principal energy level do not have the same energy (degenerate). Sublevels vary in energy within a principal quantum level.

s<p<d<f


Hydrogen vs. Polyelectronic

HISTORY OF THE PERIODIC TABLE

EARLY GREEKSEarth AirFire Water

DOBEREINER (1780-1849)

Johann Dobereiner was the first chemist to recognize patterns and found several groups of three elements that have similar properties.

chlorine, bromine and iodinecalled triads.

NEWLANDSJohn Newlands suggested that

elements should be arranged in octaves, based on the idea that certain properties seemed to repeat for every eighth element in a way similar to the musical scale.

MEYER AND MENDELEEV

The present form of the periodic table was conceived independently by two chemists: Meyer and Mendeleev. Usually Mendeleev is given most of the credit, because it was he who emphasized how useful the table could be in predicting the existence and properties of still unknown elements.

MEYER AND MENDELEEVIn 1872 when Mendeleev first

published his table, the elements gallium, scandium, and germanium were unknown. Mendeleev correctly predicted the existence and properties of these elements from the gaps in his periodic table. Mendeleev also corrected the atomic masses of several elements.

MENDELEEV’S TABLE

THE AUFBAU PRINCIPLE AND THE PERIODIC TABLE

THREE RULES FOR ORBITAL CONFIGURATIONAufbau principle – As protons are

added, so are electrons, and fill in orbitals in order of energy levels.

Pauli Exclusion – Two electrons with opposite spins can occupy an orbital.

Hund’s rule – The lowest energy configuration for an atom is the one with one unpaired electrons in each degenerate orbital. (Electrons don’t like roommates)

VALENCE ELECTRONSValence electrons are the

electrons in the outermost principal quantum level of an atom. These are the most important electrons because they are involved in bonding.

The inner electrons are known as core electrons.

VALENCE ELECTRONSThe elements in the same group

have the same valence electron configuration. Elements with the same valence electron configuration show similar chemical behavior.

TRANSITION METALS

Transition metals have electron configurations that fill in the order of 4s before 3d. Copper and Chromium have a configuration that is observed different than what is expected.

Expected: Cr: 1s22s22p63s23p64s23d4

Observed: 1s22s22p63s23p64s13d5

Expected: Cu:1s22s22p63s23p64s23d9

Observed: 1s22s22p63s23p64s13d10

TRANSITION METALS

ADDITIONAL ORBITAL RULES

The (n+1)s orbital always fills before the nd orbitals. The s orbitals fill prior to the d orbitals due to the vicinity of the nucleus.

After lanthanum, which has the configuration of [Xe] 6s25d1, a group of 14 elements called the lanthanide series, or the lanthanides occurs. This seris of elements corresponds to the filling of the seven 4f orbitals.

ADDITIONAL ORBITAL RULESAfter actinium, a group of 14

elements called the actinide series or actinides occurs.

The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence electrons for the atoms in these groups.

PERIODIC TRENDS IN ATOMIC PROPERTIES

IONIZATION ENERGY

Ionization energy is the energy required to remove an electron from a gaseous atom or ion when the atom or ion is assumed to be in its ground state:X(g) X+

(g) + e-

IONIZATION ENERGYIt is always the highest-energy

electron (the one bound least tightly) that is removed first. The first ionization energy (I1) is the energy required to remove that first electron. The second ionization energy (I2) is considerably larger.

IONIZATION ENERGYThe first electron is removed from a

neutral atom, the second from a +1 cation. The increase in positive charge binds the electrons more firmly and the ionization energy increases. The trend continues for consecutive electrons removed.

Core electrons are always held tighter than valence.

IONIZATION ENERGYFirst ionization energy increases from

left to right across a period.First ionization energy decreases in

going down a group.

IONIZATION ENERGY

ELECTRON AFFINITY

Electron Affinity is the change in energy change associated with the addition of an electron to a gaseous atom:X(g) + e- X-

(g)

ELECTRON AFFINITYIf the addition of the electron is

exothermic the corresponding value for electron affinity will carry a negative sign.

The more negative the energy, the greater the quantity of energy released.

ELECTRON AFFINITYElectron affinities generally become

more negative from left to right across a period and becomes more positive down a group.As with Ionization energy. Some

exceptions occur due to repulsions and electron configuration.

ELECTRON AFFINITY

ATOMIC RADIUS

Atomic radii are measured by the distances between atoms in chemical compounds.

Covalent atomic radii are assumed to be half the distance between atoms in covalent bonds.

For metallic atoms, the metallic radii are obtained from half the distance between metal atoms in a solid metal crystal

ATOMIC RADIUSAtomic radii decrease in going from

left to right across a period because of increasing nuclear charge and decreasing shielding.

Atomic radius increases down a group, because of the increases in the orbitals sizes associated with principal quantum numbers.

ATOMIC RADIUS

THE PROPERTIES OF A GROUP: THE ALKALI METALS

INFORMATION AND THE PERIODIC TABLEIt is the number and type of

valence electrons that primarily determine an atom’s chemistry

The organization of the period table allows the prediction of electron configuration without memorization.

INFORMATION AND THE PERIODIC TABLEGroups on the periodic table have

specialized names: Alkali metals, Alkaline earth metals, Halogens, …etc.

The most basic division of elements in the periodic table is into metals and non-metals. This division affects chemical properties.Metals tend to give up electrons and

have low ionization energies. The opposite is true for non-metals.

INFORMATION AND THE PERIODIC TABLE

INFORMATION AND THE PERIODIC TABLEMetalloids are elements along the

division line and exhibit both metallic and nonmetallic properties under certain circumstances. These elements are sometimes called semimetals.

THE ALKALI METALSLithium, sodium, potassium,

rubidium, cesium, and francium are the most chemically reactive of the metals. Hydrogen is found in group 1 but behaves as a nonmetal because its very small and the electron is bound tightly to the nucleus.

THE ALKALI METALSGoing down the group the first

ionization energy decreases and the atomic radius increases. The overall density increases due to the increase of atomic mass relative to atomic size (therefore more mass per unit volume).

THE ALKALI METALSThere is a smooth decrease in

melting point and boiling points in Group 1 that is not typical for other groups.

The most important chemical property of Group 1 is its ability to lose its valence electrons. Group 1 are very reactive.

THE ALKALI METALS Hydration energy of an ion represents the

change in energy that occurs when water molecules attach to the metal cation. The hydration energy is greatest with Li+

because it has the most charge density (charge per unit volume). This means that polar water molecules are more strongly attracted to the small Li+ ions

The order of reducing abilities in an aqueous reaction is Li > K > Na

THE ALKALI METALS

THE END

Ap chem unit 7

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