Chapter 13
Electrons in Atoms

• Democritus 400 BC
1. A Greek philosopher described matter more than 2400 years ago
2. His theory: Matter could be divided into smaller pieces only so far

3. Named the smallest particle of matter “atomos”
4. His theory was ignored for a long time.
 Aristotle
1. More popular and respected
2. His theory: 4 elements fire, air, earth and water

• Dalton’s Model 1800’s
1. Chemist
2. Performed a number of experiments that lead to the acceptance of the idea of atoms
3. His theory:
-All elements are composed of atoms. Atoms are indivisible and indestructible particles

- Atoms of the same element are exactly alike
- Atoms of different elements are different
- Compounds are formed by the joining of atoms of two or more elements
(This became of the foundations for modern chemistry)

• Thomson’s Plum Pudding Model 1897
1. A scientist
2. Provided the first hint that an atom has smaller particles.
3. His theory:
Atoms were made from a positively charged substance with negatively charged electrons scattered about

4. His experiment:
- He passed an electrical current through a gas.
- When the current passed through the gas, it gave off rays of negatively charged particles
- Discovered there were smaller particles in the atom

- Called the negatively charged
“corpuscles” now known as electrons
- Since the gas was known to be neutral, he reasoned that there must be positively charged particles in the atom
- He could never find them

 Ernest Rutherford 1908
1. Gold Foil Experiment
- He fired a stream of tiny positively charged particles at a thin sheet of gold foil
- most of the + charged particles passed right through the gold atoms in the sheet of gold foil

- Some of the + particles bounced back from the gold sheet (positive repels positive)

- The experiment explained that the gold atoms in the sheet were mostly open space.
- He concluded that an atom had a small, dense, positively charged center that repelled the positively charged particles
- He called the center of the atoms the nucleus

He reasoned that all of the atom’s + particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus.

• Niels Bohr 1913
1. His theory: proposed that the electrons were in a specific energy level
2. Electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits or energy levels are located at certain distances from the nucleus.

• The Quantum Mechanical Model
(Schrodinger)
Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space.

- It is impossible to determine the exact location of an electron. The probability of where the electron is located is based on the energy the electron has.
Electron Cloud:
 Depending on their energy they are locked into a certain area in the cloud.
 Electrons with the lowest energy are found in the energy level closest to the nucleus.

• Electrons with the highest energy are found in the outermost energy levels, further from the nucleus.

Atomic Orbitals
- regions where an e resides 95% of the time
4 shapes letter shape max e orbitals s sphere p dumbell
2 1
6 3 d 4-leaf clover 10 5 f dragonfly 14 7

 Atomic Orbitals diagrams s orbital p orbitals

d orbitals

 In nature, changes generally proceeds toward the lowest possible energy level.
 High energy systems are unstable and lose energy to become more stable.
 Electrons are arranged with lowest possible energy level (electron configurations)

Aufbau Principle
 Electrons enter orbitals of lowest energy first.
 Follow the path
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d
6p 7s 5f 6d 7p

 Examples:
Show the electron configuration for
1. sodium atomic number is 11 e = 11
1s 2 2s 2 2p 6 3s 1
2. phosphorus atomic number =15
1s 2 2s 2 2p 6 3s 2 3p 3


Rule: Place the [ preceding noble gas] then the rest of the electron configuration
Examples: Br e = 35

 Hund’s Rule
There must be one electron in each orbital of a sublevel before doubling occurs
One arrow equals one electron
So: s has 2 electrons, one orbit s _____

f p ___ ____ ____ d

Exceptions: s 1 exception one electron leaves the “s” and goes to the “d”
Nb, Cr, Mo, Tc, Ru,Rh, Cu, Ag,
Au, Pt

1. Principal Quantum number (n)
The maximum distance an electron’s orbital is from the nucleus.
n = 1, 2, 3,…..

2. Orbital quantum number (l)
The shape of an electron’s orbital l = 0, 1, 2, 3, …. (n-1) s p d f

3. Magnetic quantum number (m)
Shows how the electron’s orbital is oriented in space m = l ….0….+l s __
0

p __ __ __
-1 0 +1 d __ __ __ __ __
-2 -1 0 +1 +2 f __ __ __ __ __ __ __
-3 -2 -1 0 +1 +2 +3

4. Spin quantum number (s)
States in which direction the electron spins
Uses the right hand rule from physics s = +1/2 or -1/2

Pauli Exclusion Principle
No two electrons in the same atom can have the same set of 4 quantum numbers

Example: P
1s 2 2s 2 2p 6 3s 2 3p 3
__ __ __ __ __ __ __ __ __
-1 0 +1 -1 0 +1
* * n=2 n= 3 l= 1 l = 1 m = +1 m = 0 s = -1/2 s= +1/2

State the element whose last electron has the following quantum number’s n= 5 l = 2 m = 0 s = +1/2
5d __ __ __ __ __
-2 -1 0 +1 +2
5d 3 Ta

Energy levels
 Main areas where an electron could be
 Closest to the nucleus has lowest energy
 1 s, 2 s , 3 , 4 , 5 , 6 , 7 p

Sublevels
 The letters that stand for the shapes in the different energy levels
 Example: 3s 3p 3d
Orbitals
 determine how many electrons can be held
 Ex. 3s ___ 3p ___ ___ ____

Other questions:
1. How many occupied energy levels are in an atom of Ho (#67)
2. How many occupied sublevels?
3. How many half-filled orbitals?

Electron Dot Structure
 Keeps track of valence electrons
 valence electrons – outermost electrons
 octet rule: has eight valence electrons, stable

6
2
7 3
X
8 4
1
5
(right-left-top-bottom)
Example: Se (#34)

Behavior of electrons
 Isaac Newton 1700
1. Thought light as consisting of particles
2. Wave phenomenon

Light and Atomic Spectra
 Light is wavelike
 electromagnetic radiation: includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, xrays, and gamma rays
 waves travel in a vacuum at a speed of
3.0 x 10 10 cm/s or 3.0 x 10 8 m/s

 wavelength : (λ) the distance between the crest
 amplitude : wave’s height from the origin to crest

• Frequency: (ν) the number of wave cycles to pass a given point
 c = λν where c = speed of light
λ = wavelength
ν = frequency units: c is m/s
λ is m
ν is hertz (Hz) = s -1

• Example: What is the wavelength of the red light emitted by a barium lamp if the frequency is 3.25 x 10 14 s -1 .
 Atomic emission spectrum: the relative intensity of each frequency of electromagnetic radiation emitted by the element’s atoms or the compound’s molecule when they return to the ground state

Einstein and Planck
Photoelectric effect
1. Reflected certain colors off a piece of metal
2. Noticed electrons were released
3. Noticed not all colors did this

4. Found that certain color had specific frequencies
5. Therefore light travels as particles called photons
Planck’s constant
E = hv h= Planck’s constant
= 6.626 x 10 -34 Js

6. Light travels as both waves and particles
Summary
1. Energy caused by electrons jumping from high levels to low levels
2. Loss of energy is given to a photon of light

3. Only specific colors, frequencies, energies and jumps de Broglie’s equation predicts that all matter exhibits wavelike motions

Classical mechanics vs. Quantum mechanics
1.Classical mechanics explains the motions of objects larger than atoms. The object gains or loses energy in any amount.
2.Quantum mechanics explains the motions of subatomic particles and atom as waves.
These particles gain or lose energy in packages called quanta.

Heisenberg Principle
1. Impossible to know both the position and the path of the electron
2. Works better with smaller objects like an atom than larger objects.