Electron Configuration

Atomic Models and Energy Levels

Models of the Atom-History in Models Democritus (450BC, GR): smallest bit

that is still that element. Joseph Proust (1794 FR): proposed that

compounds form in definite proportions. Proust’s experiments and law influenced Dalton.

Dalton (1808 UK): the atom is a solid ball of the element. Know the five parts of his theory.

Faraday (1830 UK): the atom has charged pieces-magnetic field affected the stream of particles in a cathode ray tube.

JJ Thomson (1903, UK): the atom is positively charged with negative bits embedded in it (plum pudding). Cathode rays are charged particles.

Science Museum, UK

Univ. of Oregon

Atomic Models 20th Century Rutherford (1911 NZ): the gold

experiment demonstrated that most of the atom is empty space. It is a massive tiny nucleus with electrons in a cloud.

Niels Bohr (1913 DK): applied quantum theory of Einstein and others to atoms.

James Chadwick (1932 UK): proved the existence of the neutron with its slightly greater mass. Nucleus has both protons and neutrons.

BohrModel

Updated Model

This includes 1932 addition of neutrons

How to Know the Number of Electrons Subatomic particle

Atoms are composed of protons, neutrons and electrons.  The atomic number (Z) is equal to the number of

protons.  The number of protons = number of electrons for

neutral atom. The charge of the atom or ion is equal to the

number of protons – the number of electrons.  The number of electrons can be determined by

using the atomic number and charge: number of electrons (negative charges) + ion charge

Energy levels, subshells, and orbitals, and spin: Quantum Numbers Electrons are associated with energy, more

specifically quantum energy, and exemplify wave-like and particle-like characteristics.

Quantum: the magnitude of a physical phenomena takes on only discrete values

The electrons of a ground state electron are in the lowest principal energy levels possible.

Electrons in an orbital that is higher than ground state are excited electrons.

The electrons exist in specific energy levels, not a continuous range.

The principal energy level, n, which is numbered 1, 2, 3, 4, …

Excited Electrons When excited electrons drop back down to ground state, a

photon of light is emitted. The drop between two orbits has a specific wavelength or color. Each element has a unique pattern, a signature. This is used in mass spectrometry and in astronomy to

determine the elements present in a sample.

The visible spectrum of light and emission line spectra of hydrogen, neon, and iron. Note that the heavier an element is, the more spectral bands it has.

Subshells Electrons are placed in energy levels.  These energy levels are sub-divided into

subshells (labeled s, p, d or f).  The s subshells is the lowest energy and begins in

level 1.  The p subshells is higher energy and therefore

doesn’t begin until level 2.  The d is higher energy and begins in level 3 and the

f is even higher energy and begins in level 4.  The subshells are further sub-divided into orbitals (s

has 1 orbital, p has 3 orbitals, d has 5 orbitals and f has 7 orbitals). 

Each orbital can hold 2 electrons.

Valence Electrons: Element properties The location of valence electrons that are

available to participate in reactions can be predictive of properties and behavior of an element.

S orbital electrons—group 1 and group 2, the alkali metals and alkaline earth metals, have valence electrons in the s orbital. The principal quantum number “n” fills the s

orbital. The maximum is two electrons for any one s

orbital. What are the ionic charges for elements of these groups?

P orbitals The p block on the periodic table includes

groups 13 through 18, excepting Helium which does not have occupied p orbitals.

There is a maximum of six electrons filling the three orbital shapes with a maximum of two electrons each.

Think of the shape of the orbital as a map of where an electron is most likely to be, a map of probability. It is not a route that the particle flies along.

d Orbitals These are in groups 3 through 12, transition

elements. The Romans. This diverse group of metals

have variable oxidation states with the exception of Ag, Zn, and usually Cd.

Five orbital shapes 10 electrons

f Orbitals No group number. The Actinides. These larger atoms exhibit some unusual

properties (Neodymium and Praseodymium) The Lantanides. These are the heavy elements which are

unstable and radioactive. Seven orbital shapes 14 electrons

Periodic Table and Orbitals

Aufbau Principle Rules for writing electron configurations

The Aufbau principle states that energy levels must be filled from the lowest to the highest and you may not move on to the next level unless the previous level is full.  Use the periodic table as a guide (read left to right):

Aufbau Diagram (Diagonals) It is important to keep in mind

that the Aufbau principle represents an approximate trend that holds in most cases. There are however exceptions to these rules.

Why is gold, a noble metal, that is positioned in the middle of the transition elements? Exceptions tend to be where

shells can be filled (or half-filled-one electron in each orbital).

List the Levels and Sublevels in Order Start from the lowest energy level to the

highest. Use the periodic table as a guide.

Energy Levels Energy increases as electron positions move

away from the nucleus.

Hund’s Rule: or I want my own orbit. Hund’s Rule says that when placing electrons in

orbitals of equal energy, place one in each orbital before doubling up in order to arrive at the lowest energy configuration. 

The Pauli Exclusion Principle states that when electrons do share an orbital, they must be of different “spin.” No two electrons can have the same quantum numbers: each one has an unique number address.

The Elements Each time an electron

is added it goes to an open orbital in the same energy level FIRST.

In larger atoms, both the s and p may have one electron only to get a configuration with one electron in every orbital. See Cr and Cu.

This creates a more balanced geometry.

Periodic Table

Lewis Dot Table (s and p orbitals combined)

s1 s2 p1 p2 p3 p4 p5 p6

Noble: complete shells

Lewis Dot with Levels

As electrons are added, they fill electron shells in an order determined by which configuration will give the lowest possible energy. The first shell (n=1) can have only 2 electrons, so that shell is filled in helium, the first noble gas.

In the periodic table, the elements are placed in "periods" and arranged left to right in the order of filling of electrons in the outer shell. So hydrogen and helium complete the first period.

Quantum Numbers Principal quantum number (n) is the shell

(energy level) that an electron belongs to. If an electron is in the lowest possible energy state, then

it is in its ground state. If an electron is in a higher energy state, then it is in an

excited state. The total number of orbitals for n = n2. 1st level has 1*1

orbitals. 2nd level has 2*2 orbitals or 4 total. n may equal 1, 2, 3, 4, 5, 6, …

Angular momentum (secondary, azimuthal) quantum number (l) specifies the shape of the orbit or subshell. l 0 1 2 3 4 Letter s p d f g

Quantum Numbers Magnetic quantum number (ml) specifies the orientation

in space. For s orbitals, l is zero. For s orbitals, ml is zero For p orbitals, l is -1, zero, +1 (orientation to x, y, or z axis) For d orbitals, l is -2, -1, zero, +1, +2 (five orientations) For f orbitals, l is -3, -2, -1, zero, +1, +2, +3 (seven

orientations) ml = -l, …, 0, …, l

Spin quantum number (ms) specifies one of two possible spin orientations within a single orbital. This is often referred to as up or down.Pauli Exclusion Principle states that no two electrons can

have exactly the same four values for their quantum numbers. Two electrons in the same orbital must have opposite spins.

Quantum Numbers (1st Three) n l ml Number of Orbital Number of

(energy) (orientation) orbitals Name electrons (shape)

1 0 0 1 1s 2 2 0 0 1 2s 2 1 -1, 0, +1 3 2p 6 3 0 0 1

3s 2 1 -1, 0, +1 3 3p 6 2 -2, -1, 0, +1, +2 5 3d 10 4 0 0 1 4s

2 1 -1, 0, +1 3 4p 6 2 -2, -1, 0, +1,

+2 5 4d 10 3 -3, -2, -1, 0, +1, +2, +3 7 4f 14

 Spin (ms)is written as -½ or +½ .

Resources http://chemwiki.ucdavis.edu/Inorganic_Chemis

try

http://www.colorado.edu/physics/2000/quantumzone/bohr.html Bohr model animation with photon interactions.