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: nucleus very small:
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 don’t 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.
E3 n=3
n=3 E2
n=2
n=2 E1
n=1
n=1

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
2n2 n 32 4 18 3 8 2 1
Max # of Electrons
Orbit

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
don’t know absolute energy of energy levels, but
observe light emitted due to energy changes

20. ladder often used as analogy for energy levels of atom
How is this one different?
Potential Energy

21. Ground State vs. Excited 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

23. Success of Bohr’s Model
Bohr’s model could predict frequencies in emission spectrum of hydrogen
Predicted correct size of H atom
Unfortunately, didn’t work for anything with more than 1 electron

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

25. What is total # of occupied principal energy levels in atom of neon in ground state?1 2 3 4

26. What is total # of fully occupied principal energy levels in atom of nitrogen in ground state?1 2 3 4

27. What is total # of electrons in completely filled fourth principal energy level?8 10 18 32

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

29. Which electron configuration represents atom in excited state?
2-8-1
2-8
2-7-1

30. Which electron configuration represents atom of Li in an excited state?
1-1
1-2
2-1
2-2

31. The characteristic bright-line spectrum of atom is produced by its
Electrons absorbing energy
Electrons emitting energy
Protons absorbing energy
Protons emitting energy