Electron ConfigurationsStudents will understand the factors that

lead to significant advances in wave mechanics, quantum theory, and atomic

structure and function

I. Radiant Energy

A. General Information› 1. electron behavior has been studied

through light› 2. remember, light IS radiant energy› 3. originally considered to be wave energy

alone› 4. in the 1900’s, scientists determined light

behaved like a particle So it is both!

I. Radiant Energy

B. Waves1. Light waves are electromagnetic waves

• Called electromagnetic radiation (ER)• X-rays, gamma rays, and radio waves are,

also, a part of ER

2. Electromagnetic waves consist of electric and magnetic fields oscillating at right angles.a) All waves are described by 4

characteristics

Electromagnetic Radiation

I. Radiant Energy

B. Waves2. Electromagnetic waves consist (CONT.)

a. all waves are described by 4 characteristics

1. Amplitude – the height of the wave measured from its origin to its peak or crest• The brightness or intensity of light is dependent

on this part of the wave.

2. Wavelength – the distance between successive crests• The distance traveled in a full cycle • Visual light has a range between 400 to 750 nm

(10-9)

I. Radiant Energy

B. Waves2. Electromagnetic waves consist (CONT.)

a. all waves are described by 4 characteristics

3. Frequency – how fast the wave oscillates up and down (during

a given time, usually 1 second) The unit is cycle/s or hertz (Hz) 1 Hz = 1 cycle/s

Ex. FM Radio 93.1 MHz = 93.1 x 106 cycle/s Ex. Visual light is between 4 x 104 cycle/s and 7 x

1014 cycle/s

Wavelength and Amplitude

I. Radiant Energy

B. Waves2. Electromagnetic waves consist (CONT.)

a. all waves are described by 4 characteristics

4. Speed – which is a constant value• Called “the speed of light” = 3/00 x 108 M/S.• This creates a direct relationship between

wavelength and frequency.• The shorter the distance, the greater the

oscillations.• The longer the distance, the fewer the

osscilations.

I. Radiant Energy

B. Waves2. Electromagnetic waves consist (CONT.)

b. The relationship between frequency and wavelength is a mathematical expression. λ (lambda) = wavelength, V (nu) =

frequency, c = speed of light λ = v * c Ex. Helium – neon laser has a wavelength of

633 nm…v = ? 4.74 x 10-14 s-1

I. Radiant Energy C. Electromagnetic Spectrum

1. Prisms separate light into the different wavelengthsa. A rainbow is all of the light in the visible

spectrum (ROY G BIV)• Violet has the shortest wavelength, Red has

the longest wavelengthb. Visible light constitutes of a very small portion

of the electromagnetic spectrum.c. The rest of the electromagnetic spectrum is

invisible to the eye.• Consists of Gamma rays, X-rays, UV, visible, infra

red (IR), Microwaves, TV waves, and Radio• From smallest wave to longest wave

Visible Part of the Spectrum

II. Quantum Theory

A. General Information› 1. Hot objects emit electromagnetic waves

(why?) a. first emits heat (IR energy/light) b. begins to glow (Red to yellow to white for

metal) Electric stove tops

› 2. Barium and Strontium emit green and red colors (why?)

› 3. Gases give off specific colors of light when heated

II. Quantum Theory B. Planck’s Theory

1. Max Planck theorized the spectrum of radiation emitted changes with temperature

2. Theorized energy emitted or absorbed is restricted to “pieces” of particular size”

3. Proposed – There is a fundamental restriction on the amounts of energy that an object emits or absorbs, which are called quantum ( meaning fixed amount) Derived from the concept of the relationship

between frequency (v) and the energy (E) with which it is associated

Plank determined the energy constant, known as Plank’s constant (h) with a value of 6.6262 x 10-34 J/s.

II. Quantum Theory

B. Planck’s Theory4. Plank’s equation is E = H * v5. Quantum of energy of extremely small, so

it looks like a continuous climb

II. Quantum Theory

C. Photoelectric Effect1.Electrons are ejected from metal when light

is shined on it.2.A minimum frequency of light is needed to

release electron• Ex. Sodium metal wont release electrons

with red light, but will with violet light

3. Light consists of quanta of energy that behave like tiny particles of light.• Called photons• Photon energy is equal to Plank’s energy

Photoelectric Effect

Photoelectric Effect

III. Atoms: A Second Look A. Line Spectra

1. A line spectrum is a spectrum of colors created from a prism

2. Elements emit light when they are vaporized

3. Each element has a unique line spectrum• An atomic fingerprint

4. Each element when placed under a flame appears as a color• Salt – yellow, because of sodium• Lithium – red• Potassium – blue w/ red• Neon – red• Nitrogen - orange

Line Spectra for H, Ne, and Fe

III. Atoms: A Second Look

B. Bohr Model (Neils Bohr)1. Used Rutherford’s planetary model to help

explain element emissions2. Orbitals around the nucleus were based

on quanta and given a quantum number • Lowest level, n = 1 (Ground state)

3. When the electron absorbs enough energy, it will jump to the next energy level (excited state)• n = 2, 3, 4, etc.

4. Light is emitted as the electron “falls” back to the ground state

Bohr Model

III. Atoms: A Second Look

C. Matter Waves1. Louis De Broglie theorized that matter

has a dual nature.• Believed matter should have wave like

behavior and exhibit wavelengths • Called it matter waves• Came up with a mathematical formula relating

the mass and velocity of a moving particle and its possible wavelength.

• This finding was used to create the electron microscope.

• For waves to be seen from objects the mass must be very small.

III. Atoms: A Second Look

D. Heisenberg’s Uncertainty Principle1. Werner Heisenberg proposed that the

position and the momentum of a moving object cannot be simultaneously measured and known exactly. • Uncertainty Principle

2. Hard to predict where a particle will be in the future.

IV. New Approach to the Atom

A. General Information1. 3 known concepts shape the new

looka. Energy is quantizedb. Electrons exhibit wavelike behavior c. Impossible to know exactly where an

electron is in space

2. These 3 concepts lead to the quantum-mechanical model

IV. New Approach to the Atom

B. Probability and Orbitals1. General Information

a. Consider an electrons place around the nucleus as a blurry cloud

b. The cloud’s density is greater where there is a higher probability that the electron is present• Called electron density

IV. New Approach to the Atom

B. Probability and Orbitals1. General Information

c. An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found • Orbitals have characteristic shapes (not energy)• Draw orbitals based on where they are likely to

be located 90% of the time • Different orbitals are designated by different

letters• S, p, d and f• S = spherical shaped• P = dumbbell shaped• D & f = complex

Orbitals S, P, and D

IV. New Approach to the Atom

B. Probability and Orbitals2. Orbitals and Energy

a. The Principle Energy Levels in an atom are designated by the quantum number (n)• n is the principle quantum number

IV. New Approach to the Atom B. Probability and Orbitals

2. Orbitals and Energyb. The energy of the electron increases as n increases (1, 2, 3, 4, 5, 6)

• Each energy level is divided into one of more sublevels

• The number of sublevels in each principle energy level = quantum number • For example 1 = 1 sub, 2 = 2 sub, 3 = 3 sub, etc

• The sublevels are indicated by a letter. For example• n = 1; 1s• n = 2; 2s 2p• n = 3; 3s 3p 3d• N = 4; 4s 4p 4d 4f

IV. New Approach to the Atom B. Probability and Orbitals

2. Orbitals and Energyb. The energy of the electron increases as n increases (1, 2, 3, 4, 5, 6)

The number of orbitals in each sublevel is always equals the quantum number n=1; 1s : 1 spherical orbital n=2; 2s : 1 spherical orbital (larger than 1s) 2p: 3 bell shaped orbitals n=3; 3s I spherical orbital (larger)

3p: 3 bell shaped orbitals (larger than 2p) 3d: 5 complex (d) orbitals

n=4; 4s: 1 spherical orbital (larger() 4p: 3 bell shaped orbitals (larger)

4d: 5 complex orbitals (larger than 3d) 4f: 7 complex orbitals

Orbitals S, P, D, and F

IV. New Approach to the Atom B. Probability and Orbitals

3. Electron Spina. Electrons spin on their access (2 ways only)

• Can spin clockwise and counterclockwise

b. Spinning charges create magnetic fields• Clockwise in N ↑• Counterclockwise is N ↓

c. Can have parallel spins or opposite spins• If opposite spins, cancel the magnetic pull• If parallel, they create magnetic effect

d. Pauli exclusion principle• Each orbital in an atom can hold only 2 atoms with

opposite spins.

Opposite Spins w/ Mag fields

V. Electron Configurations

A. General Information1. Electron configuration is the distribution

of electrons among the orbitals of an atom

2. Electron configurations describe where the electrons are found and what energies they possess

3. Electron configurations of atoms are determined by distributing the atom’s electrons among levels, sublevels, and orbitals based on a set of stated principles.

V. Electron Configurations

B. Determining Electron Configurations1. Easy once you learn the energy levels

of the orbitals within each principle energy level, the s-sublevel is the lowest level.

2. When electrons populate the lowest energy orbitals, they are in the ground state.

Example of Electron Configuration

V. Electron Configurations

B. Determining Electron Configurations3. The electron locations can be predicted by

using the Aufbau principle, the Pauli Exclusion principle, and Hund’s Rule a. Aufbau – electrons are added one at a time, to the

lowest energy orbitals available• Until all electrons are accountable

b. Pauli Exclusion principle – An orbital can hold a maximum of 2 electrons

• To occupy the same orbital, the electrons must have opposite spins (called paired electrons)

c. Hund’s – Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons result.

Aufbau & Hund’s

Aufbau Principle

Short-cut to Electron Configurations

V. Electron Configurations B. Determining Electron Configurations

› 4. Arrows represent electrons, boxes represent orbitals called orbital diagrams (show electrons in orbitals) Ex. 6C : 1s 2s 2p

↑ = counterclockwise spin ↓ = clockwise spin

e. electron configurations are created from electron diagrams (and vice versa)

ex. 6C = 1s22s22p2

• Exponents give you the number of electrons in each energy level.

↑↓

↑↓

↑↑↓

↑

Orbital Diagram– C

Correct and Incorrect Orbital Diagrams

Orbital Diagram– P

V. Electron Configurations

C. Exceptions to Aufbau’s Principle1. Some elements don’t follow the rule

a. “They are interesting”• Ex. Cr-23, Cu-29

• Expected: Cr-23: 1s22s22p63s23p64s23d4

Actual: 1s22s22p63s23p64s13d5

• Expected: Cu-29: 1s22s22p63s23p64s23d9

Actual: 1s22s22p63s23p64s13d10

b. Cause by interactions of electrons in orbitals with very similar energies

Periodic Table and S, P, D, and F Orbitals