Chapter 4 - Electrons

Properties of Light

What is light?• A form of electromagnetic radiation:

• energy that exhibits wavelike behavior as it travels through space

Electromagnetic radiation is classified into two types:• Non-ionizing Radiation:

• Transfers energy causing vibrations, electron excitation, and heat

• Parts of the spectrum: radio, microwaves, infrared, and visible

• Ionizing Radiation:• High energy that ejects electrons and transforms

molecules into reactive unstable fragments• Parts of the spectrum: UV, X-ray, Gamma Ray

Radiation is organized on the Electromagnetic Spectrum

Wave PropertiesWavelength (λ) : • distance between

corresponding points on adjacent waves

• Unit: meters or nanometers

Frequency (ν):• number of waves that pass a

given point per unit time (usually 1 sec)

• Unit: 1/s = s-1 = Hertz (Hz)

Properties of Light• How are frequency and wavelength related?

c = λνc : speed of light (m/s)

c = 3.00 x 108 m/sλ : wavelength (m)

ν : frequency of wave (s−1 = 1/s = Hz)

Max Planck proposed …

Energy needed for electrons to move was quantized, or a quantum of energy is needed to move an electron.

E = hνE : energy of light emitted (J)

h : Planck’s constant (J·s) h = 6.626 × 10−34 J·s

ν : frequency of wave (s−1 = 1/s = Hz)

E = hn

E = hn

Electrons can only move with a quantum of energy.

Every color represents a different amount of energy released.

Louis de Broglie proposed …• that electrons be considered waves

confined to the space around an atomic nucleus.

• If electrons exist as waves, then they must have specific frequencies

• If they have specific frequencies, they also have specific energies

So energy is quantized!

ELECTRON BEHAVIOR

Who made this model of the atom?

Electrons are always moving!

• The farther the electron is from the nucleus, the more energy it has.

• Electrons can change energy levels and emit a photon of light.

• Photon: packet of energy• Ephoton = hν

Electron Movement Vocabulary

• Ground State: electrons in the lowest possible energy level

• Excited State: electrons absorb energy and move to a higher energy state

• Emission : when an electron falls to a lower energy level, a photon is released

• Absorption : energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level

• How much energy must it absorb?• Quantum: amount of energy needed to move between energy

levels

Electron Movement

• When electrons move from the excited state to the ground state they emit a photon of light

• The light emitted by an element is viewed as a bright line emission spectrum

• Each band of light represents the energy released by an electron when it moves from higher to lower energy

Line Emission Spectrum of Hydrogen Atoms

Because electrons move so much…

We cannot know an electrons exact speed and position at the same time

Heisenberg Uncertainty Principle:

Schrodinger Wave Equation

• Equation that describes the wave & particle nature of an electron

• Don’t know the exact position…• Predict where electron will be most of the time

Where 90% of thee- density is foundfor the 1s orbital

Electron Location

DO NOW• How would you describe the location of your home?

• Like giving your house an address, each electron in an atom is assigned an address or location inside the atom.

• Remember: Do electrons ever stop moving?

We do not know where the electrons are, but we can

describe where they might be

• The address of all the electrons in an atom is called an Electron Configuration

Sodium #11

1s22s22p63s1

First Location: ENERGY LEVEL

n = 1, 2, 3, 4 …

ENERGY LEVEL

SUB-LEVEL

Second Location: SUBLEVEL

These are the probability regions called: s , p , d , f

SUB-LEVEL

ORBITAL

Ene

rgy

1s

2s

3s

4s

5s

2p

3p

4p

3d

Third Location: ORBITALs = 1 p = 3 d = 5 f = 7

Sublevels on the Periodic Table

Examples:

• What is the electron configuration of Mg?

• What is the electron configuration of Cl?

How to fill in an Orbital Diagram

Ene

rgy

Step 1 – Aufbau Principle:Electrons are represented as arrows and drawn one at a time.Electrons start at the lowest possible energy.

1s

2s

3s

4s

5s

2p

3p

4p

3d

Element: HydrogenNumber of Electrons:

Electron Configuration:

Ene

rgy

Paired electrons

Step 2 – Pauli Exclusion Principle:Electrons have opposite spins.Draw one up and one down.

1s

2s

3s

4s

5s

2p

3p

4p

3d

Element: HeliumNumber of Electrons:

Electron Configuration:

Ene

rgy

Step 3 – Hund’s Rule:Electrons occupy equal energy orbitals so that a maximum number of unpaired electrons resultsDraw electrons in each orbital first then pair them.

Unpaired electrons

1s

2s

3s

4s

5s

2p

3p

4p

3d Element: OxygenNumber of Electrons:

Electron Configuration:

What is the electron configuration of S?E

nerg

y

1s

2s

3s

4s

5s

2p

3p

4p

3d

Element: SulfurNumber of Electrons:

Electron Configuration:

What is the electron configuration of Sr?E

nerg

y

1s

2s

3s

4s

5s

2p

3p

4p

3d

Element: StrontiumNumber of Electrons:

Electron Configuration:

Valence Electrons

Why is location important?

location influences behavior

Valence Electrons• valence electrons:

• electrons in the outermost energy level• can be lost, gained, or shared• electrons that are responsible for reactions

• Elements in a group have similar properties because they have the same number of valence electrons

How many valence electrons in each group?

Only the main block elements!

1

2 3 4 5 6 7

8

• All electrons under the outer most level

Inner Core Electrons