Topic 3

Electron Configurations

Line Emission for HydrogenBohr said that electrons could only travel in fixed

orbits around the nucleus…..only works for HydrogenGaseous elements emitted electromagnetic radiation

when heatedQuanta – small packets of lightGround State (release) Excited State

(absorbed)Line spectrum – distinct bands (gaseous elements)Continuous spectrum – one continuous band (fluorescent

and white light)Wavelength of line created by the electron returning to

ground state is also fixedGives a specific pattern for each element

What Bohr Proposed

1. The electron on the hydrogen atom can exist only in certain spherical orbits.

2. As the distance from the nucleus increases, the energy of an electron in that orbit increases.

3. The closest orbit (energy level) is called the ground state. Higher energy levels are called excited states.

4. When an electron falls from a higher energy level to a lower energy level, it emits a definite amount of energy that is equal to the difference in the energy of the two levels.

Bohr’s Model

Ephoton = 

energy of level nfinal -energy of level ninitial

Scientists figured that since only specific frequencies of light were emitted then the energy differences between the atoms’ energy were fixed.

This is what lead Bohr to believe that a hydrogen atom exists only in very specific energy states

These are additional lines that were discovered in the ultraviolet and infrared regions of hydrogen’s line spectrum

Wave Nature of electronsWavelength – distance between 2 repeating points on a

wave (λ)•units are nm or m (m must be used in calculations)

Frequency - # of waves that pass in a second (ν) • units are s-1 or Hz

Electromagnetic radiation are all waves that can be defined with:E = hν (h = Planck’s constant = 6.626 x 10-34 J•s

and ν = frequency)C = λν (c = speed of light = 3.00 x 108 m s-1)

E = hc/λ (E has the unit of J)mc2 = hc/λ (m = mass in kg)

Cosmic X rays ultraviolet visible infrared microwaves Radio waves

High frequency Long wavelength

Any radiation will fall somewhere in this electromagnetic spectrum

DeBroglie suggested that the electron had some wavelike characteristics

Bohr suggested that the electron was a discrete particleSchrodinger developed the idea that the electron had wave

equations that involved the electron with a particle nature• dual wave-particle nature

Heisenberg Uncertainty Principle: Heisenberg Uncertainty Principle: impossible to determine simultaneously both the position and velocity of an electron or any other particle

Orbitals give the probabilityprobability of finding an electron at a given place around the nucleus

Quantum TheoryQuantum Theory: describes mathematically the wave properties of electrons and other very small particles

Quantum NumbersPrinciple Quantum # (n)

1st shell has quantum # of 1, 2nd shell has 2Maximum # of electrons is given by 2n (where n = the

energy level)

Angular Momentum Quantum # (ℓ) aka azimuthal quantum #Subshells: s, p, d, f correspond to 0, 1, 2, 3 (0 to n-1)Letters refer to 3-D shape

s orbital = spherical shapedp orbital = dumbbell shaped and align on x, y, z axesd orbitals are 4-leaf clover shaped f orbitals have more complicated shapes

Magnetic Quantum # (mℓ)refers to orientationsEach sub shell is in orbitalsThe # of orbitals that are possible is = to twice the azimuthal

quantum # plus 1 (2ℓ + 1)Possible values are -ℓ to +ℓ including 0

Spin Quantum # (ms)Each orbital holds a max of 2 electronsPauli exclusion principle says that no one electron can have the

same set of quantum numbers, so since each orbital can hold a max of 2 electrons, they must be distinguished

Can have values of + ½ or – ½

Choice of Quantum NumbersWhen there is a choice of magnetic quantum numbers, the

lowest values are chosen 1st and + ½ is chosen before – ½

Rules for Filling OrbitalsAufbau Principle - an e- occupies the lowest orbital that

can receive it.Find how many electrons are presentLowest energy orbitals are filled 1st – 1s, 2s, etc…4s has lower energy than 3d orbitals, as is the 5s and the 4d

Hund’s Rule - orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and they must have parallel (same) spinsAll 2p orbitals have same energy (all 3d, all 4f)

Electron ConfigurationsDetermining Electron Configuration

Period # shows shellBlock shows type of orbitalAdd 1 electron until the orbital is full

Anomolies to electron configurationsCr and Cu have configurations of 4s13d5 and 4s13d10 because a half

or completely filled d shell is considered to have extra stabilityThe s and d orbitals are very close in energy so it is fairly easy for

an electron to shift between them

Noble Gas ConfigurationPhosphorus becomes [Ne]3s23p3

Write previous noble gas in square brackets and then fill in orbitals as before

You will always start with the s that has the row number of the element you are working with

Orbital diagramsArrows represent the electron (and its spin) and boxes/lines

represent orbitals

Paramagnetic and diamagneticParamagnetic species are those that are attracted by a

magnet (created by unpaired electrons present in the atom)Diamagnetic species are slightly repelled by magnets and

occur when all electrons are paired

Rydeberg EquationUsed to calculate the E changes when electrons are

promoted to higher energy levels and subsequently fall back to the lower energy levels

E = the energy associated with a particular quantum #E = -2.178 x 10-18 / n2 n is the diff. between 2 levels

Can also calculate energy released by:E = (-2.178 x 10-18 / n2) – (-2.178 x 10-18 / n2)

Where the first n2is the higher energy level and the second n2is the lower energy level

By calculating the energies for 2 quantum levels and finding the difference, one can calculate the E required to promote an electron from one to another

Released energy = + value

Energy changes during transitions are proportional to the (atomic #)2

This means that if an electron is promoted from level 1 to level 5 in a species that has less p+ in the nucleus, then the same transition for a species with more protons would be more difficult

This is because the protons in the nucleus are attracting the electron to the lower E level and more E is required to promote them

Consequently, a greater amount of E is released from the one with the larger amount of protons

D block metal ionsWhen forming metal ions, d block elements lose their outer s

electrons before any d electrons

Isoelectronic – have the same electronic configuration, as a result they must be distinguished by some other means, for example the # of protons present.