Electron Configuration and Atomic Properties

Topics:Electron Spins and MagnetismOrbital EnergyElectronic Configurations of ElementsAtomic PropertiesIons

OWLs Due on 28-November.

Electron Spin and MagnetismThere is an

additional quantum number, and one that is very important!◦Spin Quantum

Number (ms)◦Electrons can have a

spin of +1/2 or -1/2◦Since electrons are

charged particles, as they move (spin, for example), they create magnetic fields.

We’ll see how we fill electrons with spins into orbitals soon, but we need to know that there are 3 types of magnetic materials:◦Diamagnetic (non-magnetic)◦Paramagnetic (weakly magnetic)◦Ferromagnetic (strongly magnetic)

Orbital Energies (single e- species)In single electron

species (hydrogen) or even some ions, the orbitals at each energy level have the same energy.◦ Even if the orbitals are

different in physical size

This is a very simplistic model that works for few species

Orbital Energies (multiple e- species)In multiple electron

species, orbital energies are intermixed among different “n” energy levels.◦ Most common◦ Leads to electron

configurations and systematic filling for most elements and ions.

The energies change because electrons now interact with other electrons◦ Repulsive forces

Electron Configuration in AtomsThe Pauli Exclusion

Principle states, simply, that no two electrons may have the same set of quantum numbers.◦ n◦ l◦ ml

◦ ms

Additionally, atomic orbitals are filled from the lowest energy up when the atom is in the “ground” state◦ Lowest energy state

Electron Configuration in AtomsThis results in the following order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

You should know up through 4fThe total number of electrons is the

same as the atomic numberWe have as many unpaired electrons in a

specific orbital (s, p, d, f) as we can fit.◦ Hund’s Rule

And what does this mean for the Periodic Table?◦ Flashback to the previous chapter……

Let’s Build Some Electron Configurations!Remember that there are always a

few exceptions◦Focus on the rules rather than the

exceptions.Remember your filling order.

◦Fill in a box diagram◦Fill in an energy level diagram◦Write it in notation◦Use the atomic number and position of

the element on the periodic table to help!!!!

Hydrogen (ground state):◦Atomic Number 1◦# of electrons = 1◦n = 1◦l = 0 (s orbital)◦ml = 0 (one s orbital)◦ms = +1/2

1s1

Helium (ground state)◦Atomic Number 2◦# of electrons = 2◦n = 1◦l = 0 (s orbital)◦ml = 0 (one s

orbital)◦ms = +1/2 and -

1/2

1s2

Lithium (ground state)◦Atomic Number 3◦# of electrons = 3◦n = 1, 2◦l = 0,1 (s and p orbitals)◦ml = 0 (one s orbital)◦ml = -1, 0, 1 (3 p

orbitals)◦ms = +1/2 and -1/2

1s2 2s1

Carbon (ground state)◦Atomic Number 6◦# of electrons = 6◦n = 1, 2 (second period on table)◦l = 0, 1 (s and p orbitals)◦ml = 0 (one s orbital), -1,0,1 (3 p

orbitals)◦ms = +1/2 and -1/2

◦1s22s22p2

What element is this?

What is the electronic configuration for Calcium?

Shorthand Notation……Sometimes we abbreviate

electron configurations using Noble Gas Notation:◦What element is this?◦[Ar]3d104s24p5

Scandium (Sc)

Back to the Periodic Table…..Mendeleev’s periodic table

There are other types of electron configurations we need to consider (and terms)Ground (lowest) versus excited

(higher) energy state….

Inner (core) shell and valence (outer) shell electrons

Atoms versus ions (a topic for later in Chapter 7)

Ground vs. Excited StateAn excited state electronic configuration is

present when an atom (or ion) absorbs energy and an electron is promoted to a higher energy level◦This can occur even if the n energy level does not

exist in the ground state. The level is still there.◦Ground State Mg (1s22s22p63s2), 12 electrons.◦Excited State Mg (just one of many possible)

1s22s22p63s13p1

The atom must absorb energy for this to happen. When it transitions back to the ground state, that exact

amount of energy is given off Light Heat Kinetic energy transferred to another atom or molecule.

Inner versus Outer ShellInner (core) shell electrons are those in full

(“closed”) n energy levels◦Ordinarily those seen in the noble gas

configuration.◦Take Aluminum for example

1s22s22p63s23p1

[Ne] 3s23p1

[Ne] electrons represent the inner or core shell electrons

The 3s2 and 3p1 electrons are the valence or outer shell.

When ions are formed, only valence electrons are gained or lost Al3+

Periodic Trends and PropertiesEffective Nuclear Charge (Z*)

◦As atomic number increases, so does the number of protons. The nuclear charge increases, which raises the energy of orbitals surrounding the nucleus.

◦Effective nuclear charge (Z*) = Z – (# of core shell electrons)

◦ It is a relative number, used for basic comparisons.

◦ It represents the nuclear charge experienced by the highest energy, valence electrons

Atomic Size

Closely related to orbital configuration, energies and effective nuclear charge

Atoms with a greater nuclear charge “pull” electrons in closer to the nucleus and are smaller◦ Opposites attract

Covalent radius: Distance between two nuclei when two atoms are bonded together (Cl2 as an example)Metallic radius: Distance between nuclei when atoms of a metallic element are near each other in a metallic crystal (say a block of Zn)

Ionization EnergyThe energy required to remove

an electron from an element in a gaseous state.

Increases across a period because of increasing orbital energy and effective nuclear charge

Ionization energy trends

Ionization energy trends

Electron AffinityThe energy change when a

gaseous atom gains an electron

More negative values represent a greater affinity.

Ions…Filled shells (n energy levels) or orbitals,

represent the most stable electron configurations◦Electrons of opposite spin are paired◦Energy levels may be full

Ions form because the energy level of that electron configuration is particularly stable

Most ions represent elements trying to achieve noble gas electron configurations

Valence (outer) shell electrons are lost (cations) or gained (anions) to produce ions.

Cations (atom loses electrons)

Anions (atom gains electrons)

Ion sizesAnions are normally larger than

their atom “parent”, because they are gaining electrons

Cations are normally smaller than their atom “parent” because they are losing electrons

Anions are generally larger than cations