Chemistry--Unit 9: Electrons in Atoms

Lecture Notes

I. Models of the Atom

A. The Evolution of Atomic Models

1. Recall J. J. Thomson (discovered the electron)’s plum pudding model,

negatively charged electrons stuck into a lump of positively charged

material

2. Rutherford, with his discovery of the nucleus, gave the atom a small,

positively charged nucleus, electrons outside

3. Bohr placed the electrons in concentric circular paths, or orbits around the

nucleus (often referred to as planetary model); first introduction of energy

levels

B. The Quantum Mechanical Model

1. Most modern model of atomic structure

2. based largely on mathematical calculations of the probability of finding an

electron in a particular region in space (Schrodinger equation)

3. The quantum mechanical model places electron in orbitals that make up

sublevels that are found on energy levels

4. The quantum part of the model from Schrodinger’s work results from the

theory of quanta

a. a quantum of energy is the amount of energy an electron needs to move

from one energy level to another.

b. electrons cannot exist between energy levels, and on a particular level

they occupy a region rather than a “spot”

c. as the jumps from energy level to energy level occur farther from the

nucleus, the amount of energy required to make the jump decreases

d. the lowest energy arrangement, the most stable arrangement, is when the

electrons are in their original locations as close to the nucleus as possible

C. Atomic Orbitals -- see illustrations of the various orbitals and sublevels at

http://www.uky.edu/~holler/html/orbitals_2.html) shows orbital shapes of each

sublevel

http://library.thinkquest.org/15567/ie4/lessons/4.html shows the individual p

orbitals and the overall p sublevel, and the individual d orbitals and the overall d

sublevel

http://www.orbitals.com/orb/orbtable.htm

Energy

Level

Principle

quantum

number (n=)

Number of

Sublevels

Types of

Sublevels

Max # of

electrons

1 1 1 s 2

2 2 2 s, p 8

3 3 3 s, p, d 18

4 4 4 s, p, d, f 32

5 5 5 s, p, d, f, g 50

6, etc. 6, etc. 6, etc. s, p, d, f, g,

h, etc.

72, etc.

(2n2)

1. Energy levels correspond to periods on periodic table

Chemistry--Unit 9: Electrons in Atoms

Lecture Notes

2. All s sublevels hold 2 electrons, all p sublevels hold 6, d’s hold 10, f’s hold

14, etc.

3. 2 electrons per orbital, so s sublevels have 1 orbital, p’s have 3, d’s have 5,

f’s have 7, etc.

4. Every sublevel has a different amount of energy associated with it; this

energy comes from how large (how many electrons) the sublevel is and how

far it is away from the nucleus

a. lowest energy sublevels are very small and close to the nucleus

b. highest energy sublevels are very large and far away from the nucleus

II. Electron Arrangement in Atoms

A. Electron Configurations

1. Electron configuration is a method to show where all electrons in an atom or

ion are located

2. The following rules help in writing electron configurations

a. Aufbau principle: Electrons enter orbitals of lowest energy first.

b. Pauli exclusion principle: An atomic orbital may describe at most two

electrons.

c. Hund’s Rule: When electrons occupy orbitals of equal energy, one

electron enters each orbital until all the orbitals contain one electron with

parallel spins

B. Quantum Numbers

1. Electron configurations provide the address for every electron in an atom,

while quantum numbers provide the address of a particular electron in an

atom.

2. Numbers are given to energy levels, sublevels, orbitals, and spins

a. Principle quantum number (n) tells energy level; 1 for 1st, 2 for 2

nd, etc.

b. Angular momentum quantum number (l) tells sublevel; 0 for s, 1 for p, 2

for d, etc.

c. Magnetic quantum number (m) tells orbital within the sublevel

1) s has one orbital, orbital quantum number is 0

2) p has three equal energy orbitals, they are arbitrarily numbered –1, 0,

+1

3) d has five, –2, –1, 0, +1, +2

4) f has seven, –3, –2, –1, 0, +1, +2, +3

d. Spin quantum number (s or ms) is always +½ or –½ and tells the

physical spin of the electron

C. Exceptional Electron Configurations

1. There is a stability in electrons and ions to obtain filled or half-filled

sublevels

2. Because of this, many of the d and f block atoms and ions shift electrons

between sublevels of close proximity, so their electron configurations are

slightly different than expected (examples, Cr and Cu, Fe and ions)

III. Physics and the Quantum Mechanical Model

A. Light and Atomic Spectra

1. The Quantum Mechanical Model arose from the study of light

Chemistry--Unit 9: Electrons in Atoms

Lecture Notes

2. The electromagnetic spectrum is made up of visible light, as well as

microwaves, radio waves, infrared, etc.

3. This link illustrates the EM spectrum http://www.yorku.ca/eye/spectru.htm

4. Light, and all EM radiation, travels in waves at a speed (ν) of 3 × 108

meters/sec

5. The amplitude of a wave is its height from origin to crest

6. The wavelength (λ) of a wave is the distance from one point in wave to the

next identical point, or the distance between crests

7. The frequency (f) of a wave is the number or wave cycles to pass a given

point in one second, measured in s–1

, which is equal to a Hertz (Hz)

8. The speed of light, and all EM radiation is given by the equation ν = f • λ

9. In the spectrum of visible light, ROY G BIV, red has the longest wavelength

and the lowest frequency, while violet has the shortest wavelength and

highest frequency

10. Elements can also give off light

a. When electrons in a vapor sample of an element are excited by passing

high voltage electricity through them, they absorb this energy and jump

to energy levels farther from the nucleus

b. When the electrons de-excite and fall back down to their original energy

levels, this energy is released in the form of visible light

c. Because each element has a unique electron arrangement, the spectrum

produced by this visible light and often even the color of the light itself

is unique to each element

d. Spectra produced in this way are called atomic emission spectra

See this website http://jersey.uoregon.edu/vlab/elements/Elements.html

B. The Quantum Concept and the Photoelectric Effect

1. Because atomic emission spectra are not continuous, as the spectrum of

sunlight, they could not be explained by the wave theory of light

a. Wave theory would dictate all forms of light have continuous spectra

b. the individual lines seen in atomic emission spectra indicate that this and

then all forms of light must also have “packet” properties, and so began

the description of light as a particle

2. Max Planck proposed the idea of quanta in relation to light and called these

light quanta photons

3. Planck also mathematically determined that the energy of a photon can be

found from the relationship E = h • f, where h is Planck’s constant, 6.6262 ×

10–34

J•s

4. This relationship holds for all EM radiation and reveals that higher energy

radiation comes from higher frequency radiation

5. The photoelectric effect also helped in solidifying the wave-particle duality

of light

a. In the photoelectric effect, some metals eject electrons when certain

frequencies of light shine on them

b. The frequency needed to cause the photoelectric effect is called the

threshold frequency

C. An Explanation of Atomic Spectra

Chemistry--Unit 9: Electrons in Atoms

Lecture Notes

1. Quantizing energy in the form of photons helped to further explain atomic

emission spectra

2. The ground state of an electron is the lowest energy level within which an

electron can be found, the excited state of the electron is when the electron

is at some higher energy level

3. Since electrons cannot exist between energy levels, a certain amount of

energy must be released from an excited electron in order for the electron to

fall back to lower energy levels and not land in between

4. This energy was quantized in the form of photons

D. Quantum Mechanics

1. Since waves can behave as particles, it was asked by De Broglie if particles

can behave as waves

2. De Broglie derived an equation that described the wavelength of a moving

particle: λ = h / m • ν where m = mass and ν = velocity

3. De Broglie’s equation predicts that all matter exhibits wavelike motions

4. Quantum mechanics then is the study of the motion of subatomic particle

size objects

5. Another part of quantum mechanics is the Heisenberg uncertainty principle

which states that it is impossible to know exactly both the velocity and the

position of a particle at the same time (the more you know about one, the

less you know about the other)