Electrons in Atoms

1
Electrons in Atoms

• Chapter 5

2
What were early steps in development of atomic
theory?

• John Dalton Billiard Ball Theory
• Atom was indivisible
• J.J. Thomson Plum Pudding Model
• Atom was composed of smaller particles

3
Rutherford Model

• nucleus contains
• all the positive charge most of mass of atom
• nucleus very small
• only 1/10,000th of atomic diameter
• electrons occupy most of volume

4
Later Models

• Bohr Planetary Model
• Schrodinger Wave Mechanical Model

5
Problems with the Rutherford Model

• Why dont electrons crash into nucleus?
• How are electrons arranged?
• Why do different elements exhibit different
  chemical behavior?
• How is atomic emission spectra produced?

6
Atomic Emission Spectra

• gas in glass tube apply voltage across ends
• produces light
• color of light depends on gas in tube
• every element produces its own unique color

7
emission spectrum of element is set of
frequencies (or wavelengths) emitted
8
Why is emission spectra useful?

• use it to determine if given element is present
  in sample
• Neon lights

9
Emission Absorption Spectra of Elements
10
Bohr Model

• Bohr - electrons in atom can have only specific
  amounts of energy NEW idea!
• each specific amount energy is associated with
  specific orbit
• electrons restricted to these orbits
• Bohr assigned quantum number (n) to each orbit
• the smallest orbit (n 1)
• closest to nucleus
• has lowest energy
• larger the orbit, more energy it has

11
Bohr Diagram

• Shows all the electrons in orbits or shells about
  the nucleus.

n2
n1
12
Bohr Model

• energy absorbed when electron
• moves to higher orbit (farther from nucleus)
• endothermic process
• energy released when electron
• drops to lower orbit (closer to nucleus)
• exothermic process

13
energy levels get closer together the farther
away they are from nucleus
Larger orbits can hold more electrons
14
Max Capacity of Bohr Orbits
15
Electron Transitions

• If electron gains (absorbs) specific amount of
  energy
• it can be excited to move to higher energy level
• If electron loses specific amount of energy
• it drops down to lower energy level

16
Hydrogen has 1 electron, but it can make many
possible electron transitions
17
Absorption Emission

• cannot easily detect absorption of energy by
  electron
• BUT
• can easily detect emission of energy by electron
• photons (light) given off as excess energy is
  released

18
Emitted Light

• energy of emitted light (E h??
• matches difference in energy between 2 levels
• dont know absolute energy of energy levels, but
• observe light emitted due to energy changes

19
(No Transcript)
20
ladder often used as analogy for energy levels of
atom How is this one different?
Potential Energy
21
Ground State vs. Excited State

• Ground state
• lowest energy state of atom
• electrons in lowest possible energy levels
• configurations in Reference Tables are ground
  state
• Excited state
• many possible excited states for each atom
• one or more electrons excited to higher energy
  level

22
(No Transcript)
23
Success of Bohrs Model

• Bohrs model could predict frequencies in
  emission spectrum of hydrogen
• Predicted correct size of H atom
• Unfortunately, didnt work for anything with more
  than 1 electron

24
Which principal energy level of an atom contains
electron with the lowest energy?

• n1
• n2
• n3
• n4

25
What is total of occupied principal energy
levels in atom of neon in ground state?

1. 1
2. 2
3. 3
4. 4

26
What is total of fully occupied principal
energy levels in atom of nitrogen in ground state?

1. 1
2. 2
3. 3
4. 4

27
What is total of electrons in completely
filled fourth principal energy level?

1. 8
2. 10
3. 18
4. 32

28
Which atom in ground state has five electrons in
its outer level and 10 electrons in its kernel?

1. C
2. Cl
3. Si
4. P

29
Which electron configuration represents atom in
excited state?

• a) 2-8-2
• 2-8-1
• 2-8
• 2-7-1

30
Which electron configuration represents atom of
Li in an excited state?

1. 1-1
2. 1-2
3. 2-1
4. 2-2

31
The characteristic bright-line spectrum of atom
is produced by its

1. Electrons absorbing energy
2. Electrons emitting energy
3. Protons absorbing energy
4. Protons emitting energy