C H A P T E R 6

ELECTRONIC STRUCTURE OF THE ATOM

COULOMB’S LAW (POTENTIAL ENERGY FORM)

• q1 and q2 are the charges of the particles

• r is the distance between particles

• U = 0, the particles have no net interaction with each other• +U, the particles have a net repulsive interaction.• -U, the particles have a net attractive interaction.

REVIEW OF ELECTROMAGNETIC RADIATION

• c = λν• λ is wavelength (in meters)• ν is frequency (in s-1)• c is speed (in m/s) • for light, c = 3.00 x108

• E = hν• h is Planck’s Constant (6.626 x10-34 J· s)

LIGHT EMISSIONS OF SOLIDS

LIGHT EMISSION OF HYDROGEN GAS

BOHR’S HYDROGEN ATOM

• Light is absorbed or emitted from electrons transitioning between energy levels.

• Since only certain energies are observed, only certain energy levels can exist.• This is called quantization of energy levels.• Think of a ladder instead of a ramp.

ABSORPTION AND EMISSION

ENERGY LEVELS OF HYDROGEN

• Why does E get closer to 0 as n increases?

MATTER WAVES

• All matter has both particle and wave properties (wave/particle duality).

• Large objects moving slowly produce waves that are too small to observe.

• For small objects moving quickly (like electrons), wave properties are important.

𝜆=h𝑚𝑣

WAVE NATURE OF ELECTRONS

HEISENBERG UNCERTAINTY PRINCIPLE

• We can’t know both the exact location and energy of an electron.

• Since we do know the energy extremely well, we don’t know the location.

SCHRODINGER MODEL OF H

• Electrons act as standing waves• Only certain wave functions are “allowed”• Wave behavior is described by a wave function, Ψ.• Ψ2 describes the probability of finding the electron

in a certain location.• Also called the electron density

ORBITALS

• Each wave function describes the shape the electron cloud can take. These shapes are called orbitals.

• We organize orbitals by shells and subshells• Shells define size and energy (n = 1, 2, 3…)• Subshells define shape (s, p, d, f)

• Each subshell has a different number of orbitals• s = 1 orbital• p = 3 orbitals• d = 5 orbitals• f = 7 orbitals

SHAPES OF ORBITALS

ENERGIES OF ORBITALS

Single electron atoms

Multi electron atomsWhy?

ELECTRON SPIN

• Electrons exhibit a magnetic field• They don’t actually spin, but we think of them as

spinning.

• They can only spin two ways. • For sake of argument, we’ll call it up and down.

• Allowed spins: + ½ and – ½ .

SPIN AND MAGNETISM

• Paramagnetic: unpaired electrons will align in presence of a real magnet.

• Ferromagnetic (real magnets): unpaired electrons aligned in the same direction.

diamagnetic paramagnetic ferromagnetic

ELECTRON CONFIGURATIONS

• A listing of how many electrons occupy each orbital.

• 3 Rules• Aufbau Principle – Electrons fill lowest energy orbitals

first.• Pauli Exclusion Principle – Each orbital can hold two

electrons if they have opposing spins.• Hund’s Rule – Each subshell is filled in a way to give the

maximum number of unpaired electrons (maximum degeneracy).

THREE DIFFERENT NOTATIONS

• Electron Configuration• List subshells and how many

electrons they contain.• 1s22s22p63s1

• Noble Gas Notation• [Ne]3s1

• where [Ne] = 1s22s22p6.

PERIODIC BLOCKS