1. Electrons in Atoms
Chapter 5

2. Do you remember the 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 atom’s volume

4. Later Models Bohr – Planetary Model
Schrodinger – Wave Mechanical Model

5. Problems Rutherford’s Model Didn’t Address
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. Bohr Model
Bohr - electrons in atom can have only specific amounts of energy NEW idea!
each specific orbit is associated with specific amount energy
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

7. Bohr Diagram
shows all the electrons in orbits (shells) around the nucleus
E3 n=3
n=3 E2
n=2
n=2 E1
n=1
n=1

8. Bohr Model energy absorbed when electron:
moves from lower to higher energy orbit (goes farther from nucleus)
endothermic process
energy released when electron:
drops from higher to lower energy orbit (gets closer to nucleus)
exothermic process

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

10. energy levels get closer together the farther away they are from nucleus – not uniformly spaced
larger orbits can hold more electrons

11. Max Capacity of Bohr Orbits
2n2 n 32 4 18 3 8 2 1
Max # of Electrons
Orbit

12. Electron Transitions
If electron gains (absorbs) specific amount of energy
it becomes excited & can move to higher energy level
If electron loses specific amount of energy
it drops down to lower energy level
it gives off photon of light (color depends on wavelength of light given off)

13. When Matter is Heated it Gives Off Light

14. Emitted Light energy of emitted light (E = h
matches difference in energy between 2 electron levels
don’t know absolute energy of energy levels, but
observe light emitted due to energy changes

15. Example: fire works, pyrotechnics, flame test
heat energy absorbed by the metal ions excites the atoms’ electrons
absorbed energy is
eventually released in
the form of light

16. Example: light bulb
electrical energy absorbed by the filament excites the atoms’ electrons
absorbed energy is
eventually released in the
form of light

17. How this works:
electrons absorb energy, get EXCITED, and “jump” to a higher energy level
after a brief time, they “fall” back to a lower energy level, giving off a specific amount of energy (a quantum amount) in the form of a photon (colored light)

18. 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

20. Absorption & Emission
cannot easily detect absorption of energy by electron
BUT
can easily detect emission of energy by electron
SEE: photons of light given off as excess energy is released

21. THE MYSTERY OF EMISSION SPECTRUMS
there are two types:
continuous spectrum
bright line spectrum

22. Continuous Spectrum
Solids, liquids, and dense gases emit light of all wavelengths, without any gaps

23. Bright Line Spectrum
thin gases emit light of only a few wavelengths so see lines of color separated with gaps between them

24. atoms cannot emit energy continuously, rather they emit energy in precise quantities

25. Atomic Emission Spectra (AKA: bright line spectra)
apply voltage across ends of
glass tube containing gas
- light is produced
color of light depends on gas in tube
every element produces its own unique color

26. our eyes see ONE color in the gas spectrum tube, however, if we use a prism we can see that each “color” is really multiple wavelengths of different colors

27. Hydrogen has 1 electron, but it can make many possible electron transitions

28. more examples:
Hydrogen:

29. Practice Q

30. Which principal energy level of an atom contains electron(s) with the lowest energy?
answer: a

31. What is total # of occupied principal energy levels in atom of neon in ground state?
neon has 10 electrons:
1st shell: 2
2nd shell: 8
answer: b 1 2 3 4

32. nitrogen has 7 electrons: 1st shell: 2 2nd shell: 5 answer: a
What is total # of fully occupied principal energy levels in atom of nitrogen in ground state?
nitrogen has 7 electrons:
1st shell: 2
2nd shell: 5
answer: a 1 2 3 4

33. What is total # of electrons in completely filled fourth principal energy level?8 10 18 32
2n2
2(42) = 32
answer: d

34. 15 electrons total: phosphorus answer: d
Which atom in ground state has five electrons in its outer level and 10 electrons in its kernel?
15 electrons total:
phosphorus
answer: d C Cl Si P

35. Which electron configuration represents atom in excited state?
2-8-2
2-8-1
2-8
2-7-1
answer: d

36. Li has 3 electrons answer: b
Which electron configuration represents atom of Li in an excited state?
1-1
1-2
2-1
2-2
Li has 3 electrons
answer: b

37. The characteristic bright-line spectrum of atom is produced by its
electrons absorbing energy
electrons emitting energy
protons absorbing energy
protons emitting energy
answer: b