Electrons in Atoms

Ch. 13

Models of the Atom

13-1

The evolution of Atomic Models

• Dalton (1766-1844): atom indivisible• J.J. Thomson (1856-1940):

– “Plum-pudding” model – negative electrons stuck in positively charged material

• Rutherford (1871-1937):– Electrons surround dense nucleus, rest of

atom is empty space

• Bohr (1885-1962):– “Planetary model,” electrons fixed in energy

levels around nucleus

The Quantum Mechanical Model

• Quantum Mechanical Model:

– Estimates the probability of finding an electron in a certain place using the Schrodinger equation.

– “Fuzzy cloud” model; where the cloud is more dense the probability of finding the electron is high, where the cloud is less dense the probability is low.

Energy Levels• Electrons move around the nucleus

in energy levels.

• Quantum of energy = amount of E required to move to a higher level.

• When they move towards the nucleus (down a level) they release energy

• When they move away from the nucleus (up a level) they require energy.

• The farther from the nucleus the energy level is, the more energy is required to move up a level (away from nucleus).

Energy Levels

• Quantum number (n) refers to an energy level– n = 1, 2, 3, 4, …7, values increase going away

from nucleus.– Each energy level fits a certain amount of

electrons:• Level 1 = 2 electrons• Level 2 = 8 electrons• Level 3 = 18 electrons• Level 4 = 32 electrons

Sublevels + Orbitals• Within each energy level there are sublevels; the

number of sublevels is equal to the quantum number.– Ex: Energy level 4 has 4 sublevels within it.

• A sublevel is made up of atomic orbitals: s, p, d, f– Orbital s fits 2 electrons total– Orbital p fits 6 electrons total– Orbital d fits 10 electrons total– Orbital f fits 14 electrons total– s fills up first, then p, then d, then f

Energy Level (n)

Sublevel/Orbital

Electrons in each

Sublevel/Orbital

Total # of electrons in

Level

1 1s 2 2

2 2s

2p

2

6 8

3 3s

3p

3d

2

6

10 18

4 4s

4p

4d

4f

2

6

10

14 32

Atomic Orbitals

Orbital Shape # of Electrons

s Spherical 2

p Dumbbell 6

d Clover-leaf 10

f Complex 14

DRAW!

s - orbital

p - orbital

d - orbital

f - orbital

Electron Arrangement in Atoms

13-2

d-1

f-2

5d1

6d1

Label + Color on white, blank, large periodic table!

Electron Arrangement in Atoms

• Period 1 - 1s2

• Period 2 - 2s22p6

• Period 3 - 3s23p6

• Period 4 - 4s23d104p6

• Period 5 - 5s24d105p6

• Period 6 - 6s24f145d106p6

• Period 7 - 7s25f146d107p6

• 1s22s22p63s23p64s23d104p65s24d105p6

6s24f145d106p67s25f146d107p6

Electron Configuration Notation• Notation used to represent electron

configurations:

– H: 1s1

– He: 1s2

– Li: 1s2 2s1

– Be: 1s2 2s2

– B: 1s2 2s22p1

Energy level Sub level/orbital

# of electrons in sublevel/orbital

d-1

f-2

5d1

6d1

Duplicate

Recommend writing configuration in each square

You Try!

• C:– (6): 1s22s22p2

• F:– (9): 1s22s22p5

• Ne:– (10): 1s22s22p6

• Na:– (11)1s22s22p63s1

• P:– (15)1s22s22p63s23p3

• Ca:– (20)1s22s22p63s23p64s2

• Ir:– (77)1s22s22p63s23p64s2

3d104p65s24d105p66s2

4f145d7

• Cm:– (96)1s22s22p63s23p64s2

3d104p65s24d105p66s2

4f145d106p67s25f7

Write the electron configurations for the following:

• Abbreviated form: shows preceding noble gas and the configuration of only the last energy level!– Mg: 1s2 2s22p63s2

• or [Ne] 3s2

– B: 1s2 2s2 2p1 • or [He] 2s2 2p1

– Si: 1s2 2s2 2p6 3s2 3p2 • or [Ne] 3s2 3p2

– Al: 1s2 2s2 2p6 3s2 3p1 • [Ne] 3s2 3p1

– Xe: 1s2 2s2 2p6 3s23p64s23d104p65s24d105p6

• [Kr] 5s24d105p6

• What happens in the fourth period?

– After 4s2, comes 3d10, then 4p6

– Scandium (#21): 1s2 2s2 2p6 3s2 3p6 4s2 3d1

• or [Ar] 4s2 3d1

– Copper is [Ar] 4s2 3d9

– Bromine is [Ar] 4s2 3d10 4p5

• What happens in the sixth period?

– After 6s2, comes 4f14, then 5d10, then 6p6

– Tungsten (W) is [Xe] 6s24f145d4

Orbital Notation Rules1) Aufbau principle: electrons enter orbitals

of lowest energy first.

2) Pauli exclusion principle: an atomic orbital may describe at most 2 electrons.

3) Hund’s rule: one electron enters each orbital until ALL orbitals contain 1 electron with parallel spins.

Light and Atomic Spectra

13-3

(only pg. 372-375)

Electromagnetic Spectrum

• Energy in the form of electromagnetic radiation (radiant energy) travels in waves

• Waves transfer the energy from one place to another

• Ex: radio waves, TV, microwave, visible light, x-rays, gamma rays, infrared, UV

• All forms of radiant energy are part of the electromagnetic spectrum

DRAW!!!

Low energy High Energy

Wavelength + Frequency• Two main properties of electromagnetic

waves:

1) Frequency

2) Wavelength

• Wavelength is the distance between two corresponding peaks or troughs.

• Frequency is the number of wave cycles per second.

• Wavelength is inversely proportional to frequency

• wavelength frequency

DRAW!!!

Wave lengthFrequency

Wave lengthFrequency

• Higher frequency waves (short wavelength) have high energy– Ex: gamma rays, x-rays, ultraviolet rays– Ex: Violet light in visible spectrum

• Low frequency waves (long wavelength) have low energy– Ex: radio waves, microwaves, infrared (heat)

waves– Ex: Red light in visible spectrum

Light and Atomic Spectra

• Electrons absorb energy and move to higher energy states/levels

• Electrons give off that energy in the form of light when they fall back down to lower energy states, or ground state.

• ALL electromagnetic waves travel at the speed of light in a vacuum – 300 million meters per second or (3.0 x 108 m/s)

•When atoms are energized by an electric current they emit light.

•When this light is passed through a prism they produce an emission spectrum.

•Each element has its own unique atomic emission spectrum fingerprint

http://www.physics.auburn.edu/~demo/modern/7b/7b_10.jpg

http://phys.educ.ksu.edu/vqm/html/emission.html