1. AP Chemistry Unit 1: Atomic Structure
Jeff Venables
Northwestern High School

2. Cathode Ray Tube – J. J. Thomson

3. Ernest Rutherford – Gold Foil Experiment

4. Isotopes, Atomic Numbers, and Mass Numbers
Atomic number (Z)
= number of protons in the nucleus.
Mass number (A)
= total number of nucleons in the nucleus (i.e., protons and neutrons).
By convention, for element X, we write ZAX.
Isotopes have the same Z but different A.
We find Z on the periodic table.
Ions: formed by gaining (- ions) or losing (+ ions) electrons.

5. Examples Complete the following table: Symbol Protons Neutrons
Electrons
13H 12 13
35Cl- 92
146 90

6. Examples Complete the following table: Symbol Protons Neutrons
Electrons
13H 1 2
1225Mg 12 13
35Cl 17 18
238U2+ 92
146

8. Types of Spectra
Emission Spectrum – a set of colored lines produced by “downward” transitions between energy levels.
produced when electrons are excited (by electricity or flame) and then return to lower energy levels.
Absorption Spectrum – a continuous spectrum with “dark lines” missing. It is produced by “upward” transitions between energy levels.
produced when white light (or IR, UV, other) is shown through a sample. Specific colors are absorbed.
The two spectra are “complimentary.” The lines and colors involved are exactly the same.

9. Bohr Model

10. Electron Configurations
It is helpful to know how the electrons are arranged within an atom.
This explains and predicts much of the chemistry involved, such as reactions that will occur and ions that will be formed.
Energy levels –
Horizontal rows correspond to energy levels
1-7

11. Energy Sublevels (s,p,d,f)
Within each energy level are sublevels
Energy level 1 has s only.
Energy level 2 has s and p.
Energy level 3 has s, p, and d.
Energy levels 4 and above have s, p, d, and f.
Orbitals – contain up to 2 electrons each.
s sublevels have 1 orbital
p sublevels have 3 orbitals
d sublevels have 5 orbitals
f sublevels have 7 orbitals

12. Electron Configurations and the Periodic Table
The periodic table can be used as a guide for electron configurations.
The period number is the value of n.
Groups 1 and 2 have the s-orbital filled.
Groups have the p-orbitals filled.
Groups have the d-orbitals filled.
The lanthanides and actinides have the f-orbital filled.

14. Writing Electron Configurations
Examples – write electron congurations: O Al Cl Ni Pu

15. Condensed Electron Configurations
Neon completes the 2p subshell.
Sodium marks the beginning of a new row.
So, we write the condensed electron configuration for sodium as
Na: [Ne] 3s1
[Ne] represents the electron configuration of neon.
Core electrons: electrons in [Noble Gas].
Valence electrons: electrons outside of [Noble Gas].

16. Lanthanides and Actinides
From Ce onwards the 4f orbitals begin to fill.
Note: La: [Xe]6s25d14f0
Elements Ce - Lu have the 4f orbitals filled and are called lanthanides or rare earth elements.
Elements Th - Lr have the 5f orbitals filled and are called actinides.
Most actinides are not found in nature.

17. Electron Configuration Exceptions
Group 6 (Cr, Mo)
Expect s2d4
Actually s1d5
Reason: special stability of half-filled sublevels
Group 11 (Cu, Ag, Au, element 111)
Expect s2d9
Actually s1d10
Reason: special stability of filled and half-filled sublevels

18. Examples – Write condensed electron configurations:B: U:
Tl:
Es: W:
Ag:
Bi:

19. Examples – Write condensed electron configurations:
B: [He]2s22p1
U: [Rn]7s26d15f3 or [Rn]7s25f4
Tl: [Xe]6s25d104f146p1
Es: [Rn]7s26d15f10 or [Rn]7s25f11
W: [Xe]6s24f145d4
Ag: [Kr]5s14d10
Bi: [Xe]6s24f145d106p3