ATOMIC STRUCTURE

Unit 2

OVERVIEW Atomic Theory

John Dalton Law of Conservation of Mass Law of Definite Proportions Law of Multiple Proportions

Ernest Rutherford Robert Millikan J.J. Thompson

Atomic Structure Protons, neutrons, electrons Atomic number Isotopes Mass number Average atomic mass

Wave nature of light Electromagnetic Spectrum C = λv

Bohr Models Photoelectric effect

Absorption/emission E = hc/ λ

Heisenberg Uncertainty Principle

Quantum numbers Pauli Exclusion Principle Hund’s Rule Aufbau Principle

Configurations (orbital, electron, noble gas) Paramagnetism/

diamagnetism Exceptions

CHEMISTRY TIMELINE

B.C.400 B.C. Democritus and Leucippos use the term "atomos”

1500's Georg Bauer: systematic metallurgy Paracelsus: medicinal application of minerals

1600'sRobert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of elements

1700s'Georg Stahl: Phlogiston TheoryJoseph Priestly: Discovery of oxygen Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook

2000 years of Alchemy

CHEMISTRY TIMELINE1800'sJoseph Proust: The law of definite proportion (composition) John Dalton: The Atomic Theory, The law of multiple proportionsJoseph Gay-Lussac: Combining volumes of gases, existence of diatomic moleculesAmadeo Avogadro: Molar volumes of gasesJons Jakob Berzelius: Relative atomic masses, modern symbols for the elements Dmitri Mendeleyev: The periodic table J.J. Thomson: discovery of the electron Henri Becquerel: Discovery of radioactivity

1900's Robert Millikan: Charge and mass of the electron Ernest Rutherford: Existence of the nucleus, and its relative size Meitner & Fermi: Sustained nuclear fission Ernest Lawrence: The cyclotron and trans-uranium elements

THE GREEKS 400 BC Democritus

Matter consists of small particles Called them “atomos” Idea rejected by peersNo scientific proof

THE GREEKS (CONT…) Aristotle

All matter continuous4 elements = earth, water, air, and fireNo scientific proof

Idea endured for 2000 years

JOHN DALTON - 1808

School Teacher Atomic Theory

1. All matter is composed of extremely small particles called atoms. There are different kinds called elements.

2. Atoms of the same element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

3. Atoms cannot be subdivided, created, or destroyed.4. Atoms of different elements combine in simple,

whole number ratios to form chemical compounds.5. In chemical reactions, atoms are combined,

separated, or rearranged but never destroyed/created.

LAWS DERIVED FROM DALTON Law of Conservation of Mass

Total mass present before chemical reaction is same as mass after chemical reaction

2H2O 2H2 + O2 If you have 10 grams of water to start, you will get 1.12 g of hydrogen and 8.88 g of oxygen

Law of Constant Composition (definite proportions) Relative numbers and kinds of atoms are constant Water is 88.8% oxygen and 11.2% hydrogen by mass no

matter how much you have

Law of Multiple Proportions If two elements combine to form more than one

compound, the masses of the two elements are in the ratio of small whole numbers

CO2 versus CO (mass ratio is 2 to 1 for oxygen)

J.J. THOMSON British Physicist Discovered electron Cathode-ray experiment Plum pudding view of atom

THOMPSON CATHODE RAY EXPERIMENT Electric current sent through gases in glass

tube called cathode-ray tube

Surface of tube opposite the cathode glowed – caused by stream of particles

Ray traveled from cathode to anode Cathode rays deflected

by magnetic field away from negatively charged object (like a magnet) Cathode rays concluded

to have negative charge

ROBERT MILLIKAN - 1909 American Physicist Charge on each electron is same Charge of electron is -1.6022 x

10-19C Calculated mass of electron as

9.10x 10-31 kg Oil drop experiment

MILLIKAN OIL DROP EXPERIMENT Drops of oil that had

picked up extra electrons allowed to fall between two electrically charged plates

Measured how voltage on plates affected rate of fall

Calculated charges of drops then deduced charge of a single electron on the drops

ERNEST RUTHERFORD Discovered nucleus Planetary model of the atom

RUTHERFORD GOLD FOIL EXPERIMENT Bombarded thin piece gold foil

with alpha particles (positively charged particle 4 times mass of hydrogen atom) Expected to pass right through gold

foil 1 in 8000 particles deflected back

toward source “As if you fired 15-inch artillery shell

at a piece of tissue paper and it came back and hit you” Concluded most of atom is empty

space except for a very small force within atom Called positive bundle of matter the

“nucleus”

MODERN ATOMIC THEORY Atom consists of proton, neutron, and

electronProton charge = +1Neutron charge = 0 (neutral)Electron charge = -1

Protons and Neutrons located in nucleus99.9% of atom’s mass is in nucleus

Electrons located outside the nucleus

ELEMENT BLOCKS

Ag107.87

Silver47 Atomic number

Name of the element

Element Symbol

Atomic mass

ELEMENT BLOCKS Atomic Number

equal to number of protons in an atom

Element Symbol First letter always capitalized If second letter exists, it is lowercase

ISOTOPES Isotopes are atoms of the same

element having different masses due to varying numbers of neutrons.

Isotope Protons

Electrons

Neutrons

Nucleus

Hydrogen–1

(protium)

1 1 0

Hydrogen-2

(deuterium)

1 1 1

Hydrogen-3

(tritium)

1 1 2

ATOMIC MASS Atomic mass is the average of all the

naturally isotopes of that element.

Isotope Symbol Composition of the nucleus

% in nature

Carbon-12

12C 6 protons6 neutrons

98.89%

Carbon-13

13C 6 protons7 neutrons

1.11%

Carbon-14

14C 6 protons8 neutrons

<0.01%

Carbon = 12.011

MASS NUMBER Mass Number = Protons + Neutrons Not found on periodic table Isotopes have different mass numbers

(due to neutrons)

SYMBOLIZING ELEMENTS

C– 12 Atomic number

Mass numberMass number

WAVE-PARTICLE DUALITY JJ Thomson won the Nobel prize for

describing the electron as a particle His son, George Thomson won the Nobel

prize for describing the wave-like nature of the electron.

The electron is a particle!

The electron is an energy

wave!

TRAVELING WAVESMuch of what has been learned about atomic structure has come from observing the interaction of visible light and matter.

WAVE THEORY OF ELECTRON 1924 De Broglie suggested that electrons

have wave properties to account for why their energy was quantized.

He reasoned that the electron in the hydrogen atom was fixed in the space around the nucleus.

He felt that the electron would best be represented as a standing wave.

As a standing wave, each electron’s path must equal a whole number times the wavelength.

DE BROGLIE

Louis deBroglie

The electron propagates through space as an energy

wave. To understand the atom, one must understand

the behavior of electromagnetic waves.

WAVES Wavelength, l

The distance for a wave to go through a complete cycle.

AmplitudeHalf of the vertical distance from the top to

the bottom of a wave.

Frequency, nThe number of cycles that pass a point each

second.

WAVES

WAVES Longer wavelength = lower frequency = lower

energy

Shorter wavelength = higher frequency = higher energy

WAVELENGTH FREQUENCY RELATIONSHIP The SI unit of frequency (n) is the hertz,

Hz

1 Hz = 1 s-1

Wavelength and frequency are related

c = ln

c is the speed of light, 2.998 x108 m/s

PRACTICE PROBLEMThe wavelength of an argon laser's output is 488.0 nm. Calculate the frequency of this wavelength of electromagnetic radiation.

c = ln

Convert nm to m 488 nm x (1 m / 109 nm) = 4.88 x 10-7 m

Then, substitute into c = λν (4.88 x 10-7 m) (v) = 3.00 x 108 m s-1 v = 6.15 x 1014 s-1 = 6.15 x 1014 Hz

ELECTROMAGNETIC RADIATION Electromagnetic Radiation

Energy in the form of transverse magnetic and electric waves.

Electromagnetic SpectrumContains all forms of electromagnetic radiation

Visible spectrumPortion of electromagnetic spectrum that we

can see (colors)

ELECTROMAGNETIC SPECTRUM

SEPARATION OF LIGHT ‘White’ light is actually a blend of all

visible wavelengths. They can separated using a prism.

LINE SPECTRA Neils Bohr studied the spectra produced when

atoms were excited in a gas discharge tube.

LINE SPECTRA Each element produces its own set of

characteristic lines

BOHR MODEL Bohr proposed a model of how electrons

moved around the nucleus.

He wanted to explain why electrons did not fall in to the nucleus.

He also wanted to account for spectral lines being observed.

He proposed that the energy of the electron was quantized - only occurred as specific energy levels.

BOHR MODEL In the Bohr

model, electrons can only exist at specific energy levels (orbit).

Each energy level was assigned a principal quantum number, n.

Energ

y

BOHR MODEL

The Bohr model is a ‘planetary’ type model.

Each principal quantum represents a new ‘orbit’ or layer.

The nucleus is at the center of the model.

TRANSITIONSELECTRON TRANSITIONSINVOLVE JUMPS OF DEFINITE AMOUNTS OF ENERGY.

ABSORPTION EMISSION Absorption – Electromagnetic radiation is absorbed

by an atom causing electrons to jump to a higher energy state (excited state).

Emission – Energy is released by an atom as particle of light (photon) as electrons fall back to the lower energy state (ground state).

Depending on frequency of photon, different colored light may be seen

PARTICLE PROPERTIES Although electromagnetic radiation has

definite wave properties, it also exhibits particle properties.

Photoelectric effect.• First observed by Hertz and then later

explained by Einstein.• When light falls on Group IA metals, electrons

are emitted (photoelectrons).• As the light gets brighter, more electrons are

emitted. • The energy of the emitted electrons depends

on the frequency of the light.

PHOTOELECTRIC EFFECT The energy of a photon is proportional to the

frequency.

(Photon energy) E= hn

The energy is inversely proportional to the wavelength (remember c = λν so v = c/λ ).

E = hc /l

h is Plank’s constant, 6.626 x 10-34 J . Sc is the speed of light, 2.998 x108 m/s

PHOTON ENERGY EXAMPLE Determine the energy, in kJ/mol of a photon

of blue-green light with a wavelength of 486 nm.

E =

=

= 4.09 x 10-19 J

h cl

(6.626 x 10-34 J.s)(2.998 x 108 m.s-1)(4.86 x 10-7 m)

DE BROGLIE EQUATION

l = wavelength, meters h = Plank’s constant m = mass, kg v = frequency, m/s

l = hmv

DE BROGLIE EQUATION Using De Broglie’s equation, we can calculate

the wavelength of an electron.

l =6.6 x 10-34 kg m2 s-1

(9.1 x 10-31 kg)(2.2 x 106 m s-1)

The speed of an electron had already been reportedby Bohr as 2.2 x 106 m s-1.

= 3.3 x 10-10 m

l =h

mv

HEISENBERG UNCERTAINTY PRINCIPLE In order to observe an electron, one would

need to hit it with photons having a very short wavelength.

Short wavelength photons would have a high frequency and a great deal of energy.

If one were to hit an electron, it would cause the motion and the speed of the electron to change.

According to Heisenberg, it is impossible to know both the position and the speed of an object precisely.

QUANTUM MODEL Schrödinger developed an equation to

describe the behavior and energies of electrons in atoms.

His equation is similar to one used to describe electromagnetic waves.

Each electron can be described in terms of its quantum numbers.

QUANTUM NUMBERS Each electron in an atom has a unique set

of 4 quantum numbers which describe it.

Principal quantum number

Angular momentum quantum number

Magnetic quantum number

Spin quantum number

QUANTUM NUMBERS Principal quantum number, n

Tells the size of an orbital and largely determines its energy.

n = 1, 2, 3, ……

QUANTUM NUMBERS Angular momentum, l

The number of subshells that a principal level contains. It tells the shape of the orbitals.

l = n – 1 to 0 Orbitals

An orbital is a region within an energy level where there is a probability of finding an electron

Orbital shapes are defined as the surface that contains 90% of the total electron probability.

ORBITAL SHAPES

QUANTUM NUMBERS Magnetic quantum number, ml

Describes the direction that the orbital projects in space.

ml = -l to +l (all integers, including zero)

For example, if l = 2, then ml would have values of -2, -1, 0, 1 and 2.

Knowing all three numbers provide us with a picture of all of the orbitals.

QUANTUM NUMBERS Pauli added one additional quantum

number that would allow only two electrons to be in an orbital.

Spin quantum number, ms. It can have values of +1/2 and -1/2

Pauli exclusion principlePauli proposed that no two electrons in an

atom can have the same set of four quantum numbers

QUANTUM NUMBERS

OTHER RULES Aufbau Principle

Electrons are placed into orbitals, subshells, and shells in order of increasing energy

OTHER RULES Hund’s Rule

The most stable arrangement of electrons in a subshell is the one in which electrons have the most number of parallel spins possible.

ORBITAL NOTATION OF ELECTRONS Graphical representation of an

electron configuration One arrow represents one electron Shows spin and which orbital within a

sublevel Follow all rules(Aufbau principle, two

electrons in each orbital, etc. etc.)

ORBITAL NOTATION Use atomic number as number of

electrons in an atom

He

Be

Mg

Si

Ne

MAGNETISM Diamagnetism

Elements have all of their electrons spin paired

All of an element’s subshells are completedNot affected by magnetic fields

ParamagnetismNot all electrons are spin paired in an

elementMost elements are thisAffected by magnetic fields

ELECTRON CONFIGURATION A list of all the electrons in an atom (or

ion)Must go in order (Aufbau principle)2 electrons per orbital, maximum

We need electron configurations so that we can determine the number of electrons in the outermost energy level.These are called valence electrons.

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

ELECTRON CONFIGURATION

2p4

Energy Level

Sublevel

Number of electrons in the sublevel

EXAMPLESHe, 2: 1s2

Ne, 10: 1s2 2s2 2p6

Ar, 18: 1s2 2s2 2p6 3s2 3p6

Kr, 36: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

PERIODIC TABLE Orbitals grouped in s, p, d, and f orbitals

(sharp, proximal, diffuse, and fundamental)

s orbitals

p orbitalsd orbitals

f orbitals

WHY ARE D AND F ORBITALS ALWAYS IN LOWER ENERGY LEVELS d and f orbitals require LARGE amounts of

energy

It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy

NOBLE GAS NOTATION A way of abbreviating long electron configurations Since we are only concerned about the outermost

electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].

Step 2: Find where to resume by finding the next energy level.

Step 3: Resume the configuration until it’s finished.

Example: [Ne] 3s2 3p5

EXCEPTIONS Remember d and f orbitals require LARGE

amounts of energy

If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)

There are many exceptions, but the most common ones are

For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.

EXCEPTIONS d4 is one electron short of being HALF

full In order to become more stable (require

less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.For example: Cr = [Ar] 4s2 3d4

Since this ends exactly with a d4 it is an exception to the rule. Thus, Cr = [Ar] 4s1 3d5

Remember, half full is good… and when an s loses 1, it too becomes half full!

d9 works the same way