1. Academic Chemistry Final Review

2. Final Exam 51 Multiple Choice Periodic Trends – definition
Atomic/Ionic Radius, Electronegativity, ionization energy, electron affinity
Accuracy/Precision
General Equations
Valence Electrons
Octet
Types of Energy
metals/nonmetals/metalloids
Pure substances – elements/compounds
Mixtures – homogeneous/heterogeneous

3. Final Exam 51 Multiple Choice Balanced Equations Stoichiometry
Polarity – bond & molecule
(s), (l), (g), (aq)
VSEPR
anion/cation
empirical/molecular formulas
Heisenburg/DeBroglie/Hund/
Temperature scales based on
family names

4. Final Exam 51 Multiple Choice covalent/ionic Rutherford
frequency/wavelength/amplitude
Effect of changes on gases
Gas pressure
parts of the scientific method
separation methods
alpha/beta/gamma
electrons on a sublevel
indicators of a chemical change

5. Final Exam 51 Multiple Choice conservation of mass/energy
significant figures
physical/chemical changes
periodic table arrangement
limiting reactant
molar mass
atom/molecule/formula unit
protons/neutrons/electrons
Thompson

6. Final Exam 9 problems limiting reactant problem
empirical and molecular formula
write a balanced formula & determine type
Differences between solids/liquids/gases
% yield
density
atom/ion/isotope
ideal gas law
names/formulas

7. Books
Books must be turned in before you leave! You may turn them in during class any day this week!

8. Review Chapters 1-4 Steps
4. hypothesis – educated guess about a problem
theory – tested hypothesis
law – tested many times, scientific community believes it to be true.
68.59m = 4 s.f.
= 4 s.f
Steps
1. Identify the problem or question
2. Gather information and observations
3. Form a hypothesis
4. Test the hypothesis
5. Analyze the data and form a conclusion

9. Review Chapters 1-4 5. 6. 23°C + 273 = 296 K °C = (62-32) °C=16.7
9. Kinetic – motion (rolling ball)
Potential – above ground level (boulder on a cliff)
Radiant – light (sun)
10. Mass – amount of material
Weight – gravity’s effect on the mass 5. 6.
23°C = 296 K
°C = (62-32) °C=16.7
16.7°C+273 = 289.7K

10. Review Chapters 1-4 11. a. element – atom b. Covalent - molecule
15. Aufbau Principle – Electrons are added one at a time to the lowest energy level available.
Pauli Exclusion Principle – An orbital can hold a maximum of 2 electrons. These electrons must have opposite spins (clockwise & counter-clockwise).
Hund’s Rule – Electrons fill orbitals so that a maximum number of unpaired electrons exist. (Ladies first!)
11. a. element – atom
b. Covalent - molecule
c. Ionic – Formula Unit
12. Mendeleev
13. Protons = 92
Neutrons (233-92) = 141
Electrons = 92
14. Mass of Carbon - 12

11. Review Chapters 1-4 16. 1s22s22p63s23p64s23d104p6 5s24d105p66s24f145d9
17. [Xe]6s24f7
18. Ion – different # of electrons giving overall charge
Isotope – different # of neutrons (different mass)
19. a. Dalton – Dalton’s atomic theory (from ideas of Democritus)
b. Thomson – Cathode Ray Tube experiments
c. Rutherford – Gold Foil Experiment
20. electron < proton = neutron < nucleus

12. “d” “p” “s” “f” Review Chapters 5-9 1. =metaloid (semi-metal)
Alkali metals
Noble gases
=metaloid (semi-metal)
Alkaline earth metals
halogens
Metals
Non-Metals 1.
Transition Metals
“d”
“p”
“s”
Inner
Transition Metals
“f”

13. Review Chapters 5-9 2. a. Cl-1 d. Al+3 b. Na+1 e. N-3 c. Mg+2 f. S-2
An octet
Ionic – metal/non-metal
Covalent – 2 non-metals
5. Valence Shell Electron Pair Repulsion Theory
6. Polar – unequal sharing of electrons – causing partial positive and partial negative
determine by difference in electronegativity

14. Review Chapters 5-9 7. Base geometry - linear Shape – linear Non-polar
Base geometry – tetrahedral
Shape – Bent
Polar
Shape – tetrahedral

15. 11. light/heat. Precipitate. Gas. Color change.
Review Chapters 5-9
8. a. sulfate c. phosphate
b. ammonium d. carbonate
9. a. synthesis d. decomposition
b. single replacement e. double replacement
c. combustion
A molecule that exists in nature as two atoms bonded together.
Br I N Cl H O F
11. light/heat. Precipitate. Gas. Color change.

16. Review Chapters 5-9 12. 3MgCl2 + Al2(SO4)3  3MgSO4 + 2AlCl3
13. a. barium hydroxide d. Ammonium hydroxide
b. iron (III) chloride e. Lead(II) sulfate
c. dinitrogen monoxide f. Copper (I) sulfide
14. a. NH4F d. Na2CO3
b. CaSO4 e. FeSO3
c. Mg3(PO4)2 f. Cr3P2
15. 2C2H O2  4CO2 + 6H2O

17. Review Chapters 10-13 a. 40.0 g c. 183.4 g b. 342 g d. 149.0 g
Empirical = lowest ratio Molecular = actual ratios
No, we cancel in criss cross to get lowest ratio
6.02 x number of particles in a mole

18. Review Chapters 10-13 5. Standard Temperature and Pressure
0°C and 1atm

19. Review Chapters 10-13 Na – 23.o x 3 = 69.0 /164.0 x 100 = 42.1%
P – 31.0 x 1 = /164.0 x 100 = 18.9%
O – 16.0 x 4 = 64.0 /164.0 x 100 = 39.0%
164.0 8.

20. Review Chapters 10-13 9.
10.
11.

21. Review Chapters 10-13 Exo – releases heat Endo – absorbs heat
13. Enthalpy change – tells us how much heat is gained or lost during a reaction
The amount of heat needed to raise the temperature of 1 gram of a substance 1 degree C
q= mC∆T

22. Review Chapters 10-13 a. P1V1 = P2V2 d. P1V1T2 = P2V2T1
b. T1V2 = T2V1 e. PT = P1 + P2 + P3…
c. f. PV = nRT
P = pressure, V = volume, T = temperature, n = moles, R = ideal gas constant ( g atm/ mol K)
The movement of one gas through another
17.