1. Electrons in Atoms Chapter 5

2. Objectives
Be able to discuss how the Bohr model of the atom was able to explain the existence of atomic emission spectra.
Be able to explain how emission spectra are formed.
Be able to explain why the emission spectrum of each element is unique.

3. Models of the Atom John Dalton (1803): solid sphere - - - -
J.J. Thomson (1897): plum-pudding - - - -
Ernest Rutherford (1911): dense nucleus +
Niels Bohr (1913): energy levels +
Erwin Schröedinger (1926): electron clouds

4. Emission Spectra
Each element had a unique “emission spectrum,” but what causes it?

5. Observing of Emission Spectra
What we see
with our eyes… =
A prism or diffraction
grating inside a
spectroscope separates
the individual colors.

6. The Bohr Model
Niels Bohr (1913): Electrons orbit the nucleus at specific energy levels; colors occur when
electrons change levels.
ultraviolet, UV
infrared, IR
higher
energy
visible
light
lower
energy

7. Production of Emission Spectra
photon
(1) A quantum (small, exact amount) of energy is absorbed by the atom.
(2) An electron in its ground state…
(3) … absorbs the energy and makes a quantum leap to an excited state.
(4) The excited electron soon drops back down…
(5) … releasing energy as a photon (particle of light).

8. Production of Emission Spectra
Each color depends on the energy difference (distance) between the levels:
small = red, orange,
medium = green, yellow
large = blue, violet.
very small = infrared (IR) very large = ultraviolet (UV)
Each atom has several energy levels, so it can produce several colors in its emission spectrum

9. Production of Emission Spectra
Each element has a unique spacing of energy levels, so it will produce a unique emission spectrum.

10. Hydrogen’s Spectrum
Bohr could only calculate the colors for H, not other “complex” elements
We’ll do the calculations in Physics next year.

11. Objectives
Understand the Heisenberg Uncertainty Principle and how it applies to atomic structure.
Be able to discuss the concepts of the atomic sublevels and orbitals.

12. Energy Levels and Sublevels
Bohr’s model was too simple, more levels must exist
orbits can’t be circular

13. Heisenberg Uncertainty Principle
Heisenberg uncertainty principle: it is not possible to precisely determine an electron’s orbital path.
observing the path alters the path
The photon changes
the path of the electron.
Werner Heisenberg

14. Electrons and Probability
90%
Erwin Schröedinger: determined the shapes of orbitals (where an e- is 90% of the time).

15. Quantum Mechanical Model
Electrons exist in orbitals of various shapes (s, p, d, f):

16. Energy Levels, Sublevels, and Orbitals4f 4d
Energy Levels, Sublevels, and Orbitals
n = 4 4p 4s 3d
n = 3 3p
sublevels (rows) 3s 2p
n = 2 2s
n = 1 1s
orbitals (represented by boxes)

18. Objectives Understand how electrons fill orbitals in atoms.
Be able to write a complete electron configuration for an element.

19. Orbital Filling Rules Aufbau principle: fill upwards from bottom
Pauli exclusion principle: up to 2 e- per orbital, must
have opposite spin (↑ or ↓) due to magnetism
Hund’s rule: place one e- into each orbital
before putting in the second e-
What are the electron configurations of H and He? 2p 2p 2s 2s
H = 1s1
He = 1s2 1s 1s ↑ ↑↓

20. Orbital Filling Rules What are the electron4p
What are the electron
configurations of O, P, Ti,
and Se? 3d 4s 3p 3s 2p 2s 1s

21. Orbital Filling Rules
What is the electron configuration of silicon and
what would the atom look like? 3d 3p ↑ ↑ 3s ↑↓ 2p ↑↓ ↑↓ ↑↓ 2s ↑↓
Si = 1s2 2s2 2p6 3s2 3p2
This is why we
don’t draw them! 1s ↑↓

22. Objectives
Be able to write a shorthand “noble gas notation” electron configuration.
Be able to use the periodic table to write an electron configuration.

23. Noble Gas Notation A shorthand noble gas notation can be used to
represent an electron configuration. 3d ↑ ↑ 3p ↑ 3s ↑↓ 2p ↑↓ ↑↓ ↑↓
same as neon (Ne) 2s ↑↓
P = 1s2 2s2 2p6 3s2 3p3 1s ↑↓
P = [Ne] 3s2 3p3

24. Overlapping Energy Levels
6d □ □ □ □ □
5f □ □ □ □ □ □ □
7s □
6p □ □ □
5d □ □ □ □ □
4f □ □ □ □ □ □ □
6s □
5p □ □ □
4d □ □ □ □ □
5s □
4p □ □ □
3d □ □ □ □ □
4s □
3p □ □ □
3s □
2p □ □ □
2s □
1s □
energy levels and sublevels overlap
the periodic table
shows this!

25. Sublevels and the Periodic Table4f 5f

26. Writing Electron Configurations3d 4p 5s 4d
What is the electron configuration for Se?
Se = [Ar] 4s2 3d10 4p4
What is the electron configuration for Tc?
Se = [Kr] 5s2 4d5

27. Objectives
Be able to explain why some elements have exceptional electron configurations.
Be able to write the most likely electron configuration of an exceptional element like silver or gold.

28. Exceptional Electron Configurations
Cr and Ag are
less corrosive
due to this stability!
Half-filled and/or completely-filled
sublevels are more stable
expected: Cr = [Ar] 4s2 3d4
actual: Cr = [Ar] 4s1 3d5
expected: Ag = [Kr] 5s2 4d9
actual: Ag = [Kr] 5s1 4d10
Usually, a single electron gets placed in a d-sublevel
before the following electrons get placed in the f-sublevel.
actual: U = [Rn] 7s2 6d1 5f 3