light and electrons. electromagnetic radiation light is electromagnetic radiation: combined electric...

Light and Electrons

Electromagnetic Radiation

Light is electromagnetic radiation: combined electric and magnetic wavesSourc

e Electric vector

Magnetic vector

direction of propagation

Electromagnetic Radiation

Light is more than what we can see…

Electromagnetic Radiation

Subatomic particles (electron, photon, proton, etc) exhibit both PARTICLE and WAVE properties. This is known as Wave-Particle Duality.

Diffraction: wave-like

Photoelectric Effect: particle-like

Electromagnetic Radiation

Wave Properties of Light:

1. It’s fast! …c = 3.0 x 108 m/s

2. It relfects, refracts, diffracts (Transverse wave)

3.

All light waves have frequencywavelength

symbol: f (Greek “lambda”)

units: “cycles per sec” = Hertz “distance” (m, nm)

where c = velocity of light = 3.00 x 108 m/sec

Electromagnetic Radiation

c = f

Increasing frequency

Electromagnetic Radiation

Example: Red light has = 700 nm. Calculate the frequency, f.

=3.00 x 10

8 m/s

7.00 x 10-7

m

4.29 x 1014

Hz f = C

Electromagnetic Radiation

Particle Properties of Light:

1. A particle of light is called a photon

2. Energy of a photon is calculated by E = h ·f

where E = energy (Joules, J)

f = frequency (Hertz, Hz, 1/sec)

h = Planck’s constant 6.63 x 10 J·s

-34

Electromagnetic Radiation

Albert Einstein postulates the Photoelectric Effect to explain two observations:

1. No electrons are observed until a minimum energy is applied.

2. Number of electrons ejected depends upon light intensity – not light frequency!

Light is created by the Photoelectric Effect

Electromagnetic Radiation

The photoelectric effect and the idea of discrete, quantized energies neatly explain the observation of emission spectra.

Electromagnetic Radiation

Example: Red light has = 700 nm. Calculate the energy per photon.

E = hf and c = f

So f = c/ and E = hc/

E = (6.63 x 10 Js)(3.0 x 10 m/s) 700 x 10 m

E = 2.84 x 10 J

-34 8

-19

-9

Electron Orbitals

While thinking about the emission spectrum of hydrogen, Neils Bohr came up with the planetary model of the atom. In this model, electrons can only orbit the nucleus at discrete distances and particular orbital shape.

Orbital model of Na

Sharp-line spectrum of H

Neils Bohr

Electron Orbitals

Electron Orbitals (n)

n = energy level or shell (n = 1, 2, 3, 4, 5, 6, 7)

1. Energy levels are whole numbers

2. The maximum number of electrons in each energy level equals 2n2.

3. The rows of the periodic table correspond to energy levels.1. Whole number energy levels – like a standing

wave

3. The rows (periods) of the periodic table correspond to energy levels.

Electron Orbitals (n)

Electron Orbitals (l)

l = subshell (s, p, d, f, g, h, i, j…)

1. s, p, d, and f are named after the four lines in the hydrogen emission spectrum…Sharp, Principle, Diffuse, Fundamental.

2. Each subshell has a different shape

3. The number of subshells in an energy level is equal to the number of the energy level.

Energy Level

Number of Sublevels

Name of sublevels

1 1 s

2 2 s, p

3 3 s, p ,d

4 4 s, p, d, f

Electron Orbitals (l)

1. Sharp, Principle, Diffuse, and Fundamental refer to the way the spectral lines look. It was thought that electrons traveling between certain energy sublevels

produced those certain lines. This was not correct, but the names stuck.

Electron Orbitals (l)

2. Each subshell has a different shape

s-orbital1.Has a

spherical shape

2.Can hold up to 2 electrons

3.Lowest energy subshell

p-orbitals

Electron Orbitals (l)

1.Said to have a “dumbbell shape”

2.Can hold up to 6 electrons

Electron Orbitals (l)

d-orbitals

1.Said to have a “clover leaf” shape

2.Can hold up to 10 electrons

combined orbitals

d-orbitals

Electron Orbitals (l)

f-orbitals 1. Can hold up to 14 electrons

combined orbitals

f-orbitals

Electron Orbitals

To write a ground-state electron configuration:

1. Determine how many electrons are present.

2. Follow the Aufbau Diagram (Diagonal Rule)

Aufbau Diagram

Electron Orbitals

Example: Write the ground-state electron configuration for nitrogen.

1. Nitrogen has 7 electrons

2. Follow the Aufbau Diagram

3. N: 1s22s22p3

Electron Orbitals

So why does it work like this?

1. Pauli Exclusion Principle – states that “no two electrons in an atom can have the same set of four quantum numbers.” In other words, no atomic orbital can contain more than two electrons.

2. Hund’s Rule – The most stable arrangement of electrons around an atom is one with the maximum number of unpaired electrons. This minimizes electron-electron repulsion.

Electron Orbitals

So why does it work like this? (cont.)

3. Aufbau Principle – Electrons occupy the lowest energy state possible.

4. Heisenberg Uncertainty Principle – The orbitals are probabilities – not shapes in space like planetary orbits. The uncertainty principle states that you cannot know the location and velocity of an electron simultaneously.

s-orbitals in Zinc

p-orbitals in Zinc

Electron Configuration Shortcut…

Electron Orbitals

Electron orbital notation goes one step further than electron configuration. It describes, specifically, each electron.

Compare them

Electron Configuration of Oxygen: 1s22s22p4

Electron Orbital Notation of Oxygen:

. . . . .

1s 2s 2p

Electron Orbitals

Orbital Notation

s . or . or . 1s 2s 3s

p . . . 2p

d . . . . .3d

f . . . . . . .4f

Electron Orbitals

Example: What is the electron orbital notation for sulfur?

. . . . . . . . .

1s 2s 2p 3s 3p

Example: What is the non-core electron orbital notation for gold?

[Xe] . . . . . .

6s 5d

Electron OrbitalsElectron Orbitals

Example: What is the non-core electron orbital notation for gold?

[Xe] . . . . . .

6s 5d

…or more likely,

[Xe] . . . . . . 6s 5d

Electrons are more stable in full or half-full orbitals.

Electron Orbitals

Octet Rule: Atoms will gain or lose electrons to achieve a full valence shell (usually this means 8 electrons).

Oxidation State: The value of the charge on an ion (positive or negative), after the atom has achieved a full valence shell.

- metals tend to lose electrons, forming positive (+) ions (cations).

- non-metals tend to gain electrons, forming negative (-) ions

(anions).

Electron OrbitalsPeriodic Table of Oxidation States