1. Chemistry Final Review
2nd Semester,

2. Unit 1: Basics How many sig figs are in the following numbers? 2300 2
314 3
2.06
0.0025 2
9.001 4
6.02x1023 3
587.0

3. Unit 1: Basics Discuss the following in terms of accuracy and precision

4. Unit 1: Basics
List the three mole conversions you learned at the beginning of the year.
1 mol = 6.02x1023 particles (atoms, molecules, ions, etc)
1 mol = molar mass in grams from the periodic table
1 mol = 22.4 L of GAS at STP  see how that makes more sense now ?

5. Unit 1: Basics Write the following in standard or scientific notation
2.71x104
27100
6.4x10-3
0.0064
2687
2.687x103
1.2654x10-2

6. Unit 1: Basics What is the equation for density? Density = mass volume
What are the common units for density?
g/mL or g/cm3

7. Unit 1: Basics
Give an example of an element and an example of a compound.
Element: He, C, Mn
Compound: CO2, NaCl, etc (anything with more than one element)

8. Unit 1: Basics Write the equation for percent yield.
% yield = actual x 100
expected
Write the equation for percent error.
% error = (actual-expected) x 100

9. Unit 2: Atoms & Periodic Table
Complete the following table:
Element/ion
Atomic number
Atomic mass
Protons
Neutrons
electrons Fe
Cl- K+ 26
55.85 26 30 26 17 18 17
35.45 18 19
39.10 19 20 18

10. Unit 2: Atoms & Periodic Table
Positive ions form when:
Atoms lose electrons (usually metals, on left of table)
Negative ions form when:
Atoms gain electrons (usually nonmetals, on right of table)
Why do atoms form ions?
To become more stable, get the configuration of a noble gas.

11. Unit 2: Atoms & Periodic Table
The halogens make a charge of ____ when they become ions. -1
The alkali metals make a charge of ___ when they become ions. +1
The alkali earth metals make a charge of ___ when they become ions. +2

12. Unit 2: Atoms & Periodic Table
The halogens make a charge of ____ when they become ions. -1
The alkali metals make a charge of ___ when they become ions. +1
The alkali earth metals make a charge of ___ when they become ions. +2

13. Unit 2: Atoms & Periodic Table
What are the three types of nuclear decay?
Alpha, Beta, Gamma
What type of particle does each emit?
Alpha = helium nucleus (2 protons, mass of 4)
Beta = electron (no mass, -1 proton)
Gamma = high energy (no mass, no proton change, but product is more stable)

14. Unit 2: Atoms & Periodic Table
Complete the following:
Type of Decay
________ 99m43Tc  9943Tc + ______
________ Am  0-1e + _____
________ Np  42He + ____
Gamma
00γ
Beta
24796Cm
17191Pa
Alpha

15. Unit 2: Atoms & Periodic Table
Draw a Bohr model for Beryllium.
Draw a Bohr model for Silicon.

16. Unit 2: Atoms & Periodic Table
There will be no electron configuration on the final (s, p, d, f).

17. Unit 3: Compounds & Bonding
Which types of elements participate in ionic bonding?
Metals (+) and non-metals (-) (also polyatomic ions)
Which types of elements participate in covalent bonding?
Non-metals and non-metals (they share electrons instead of charges sticking together)

18. Unit 3: Compounds & Bonding
Name the following compounds:
MgO
Magnesium oxide
AlF3
Aluminum fluoride
NiSO4
Nickel (II) sulfate
FeCl2
Iron (II) chloride
N2O5
Dinitrogen pentoxide
SF4
Sulfur tetrafluoride

19. Unit 3: Compounds & Bonding
Describe a polar bond.
Covalent bond (non-metal and non-metal) in which electrons are shared UNEVENLY (one atom is more electronegative than the other).
Draw a water molecule and show its polarity.

20. Unit 3: Compounds & Bonding
There will be no molecular shapes on the final.

21. Unit 3: Compounds & Bonding
List and describe the 4 types of Intermolecular forces (IMF).
Dispersion forces (weakest, between nonpolar molecules)
Dipole-dipole interaction (stronger, between polar molecules)
Hydrogen-bonding (STRONG, between molecules with N, O, or F bonded to H, explains water’s high boiling point)
Ion-molecule interaction (strongest, how water dissolves salt – pulls apart the ions)

22. Unit 4: Energy Transfer
Draw a heating curve for water.

23. Unit 4: Energy Transfer Label the Q-equations for each section. Q=mHv
Q=mcΔT
Q=mHf

24. Unit 4: Energy Transfer
Find the value for Hf, c, and Hv in your data book.
Hf = cal/g
C = 1.0 cal/g°C
Hv = cal/g

25. Unit 7: Solutions Use Table E to determine the solubility of each substance: ammonium chloride barium carbonate silver iodide mercury (II) bromide

26. ammonium chloride soluble barium carbonate nearly insoluble silver iodide nearly insoluble mercury (II) bromide slightly soluble

27. Unit 7: Solutions Use Table D: 1. How many grams of sodium nitrate will dissolve in 100 g of water at 25 C? 2. How many grams of ammonia (NH3) will dissolve in 100 g of water at 100C? 3. If 140 g of KI is dissolved in 100 g of water at 30 C, is the solution saturated, supersaturated, or unsaturated?

28. 92 g 7 g unsaturated

29. Unit 8: Equilibrium Describe a system at equilibrium for each case: ·
Unit 8: Equilibrium Describe a system at equilibrium for each case: · liquid and gas phases of the same substance in a closed container · reactants and products in a chemical system · a saturated solution

30. 33A. · liquid + gas in a closed container with equal # of molecules condensing as are evaporating; rate of condensing = rate of evaporating · rate of forward reaction is equal to rate of reverse reaction · rate of dissolving is equal to rate of precipitation or crystal growth/formation

31. Unit 8: Equilibrium Use LeChatelier’s Principle and the following reaction: PCl3(g)+3NH3(g)P(NH2)3(g)+3HCl(g) If [PCl3] increases then [HCl]… If pressure increases then [NH3]… If [P(NH2)3] increases then [HCl]…

32. 35A. If [PCl3] is increased then [HCl]… increases.
If pressure is increased then [NH3]…is constant.
If [P(NH2)3] is increased then [HCl]…decreases.

33. The last seven questions will give you practice doing basic chemistry conversions and doing stoichiometry.

34. How many moles are in 3.6 g of sodium chloride?
44.
How many moles are in 3.6 g of sodium chloride?

35. 44A mol

36. What is the volume of 2.3 moles of oxygen? (at STP)
45.
What is the volume of 2.3 moles of oxygen? (at STP)

37. 45A L

38. 46. How many atoms of mercury are in 5.0 moles of mercury?

39. 46A x 1024 atoms

40. 47. How much space does 64 g of oxygen occupy?

41. 47A L O2

42. 48. What is the mass of 3.75 x 1022 molecules of CO2?

43. 48A g CO2

44. 49. In ammonia production, nitrogen and hydrogen are synthesized into ammonia (NH3). What mass of ammonia will be produced from 1.5 kg of nitrogen assuming that hydrogen is in excess? (Hint: first write the complete, balanced equation)

45. 49A. N2 (g) + 3H2 (g)  2NH3(g) 1.8 kg NH3

46. 50. A solution made from 5.0 g of copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed.
a) Write the complete, balanced equation for the reaction.
b) Write the complete ionic equation for the reaction.
c) Write the net ionic equation for the reaction.

47. 50. A solution made from 5.0 g of copper (II) sulfate is mixed with a solution containing excess calcium nitrate. A precipitate of calcium sulfate is formed.
Continued:
d) How much (mass) calcium sulfate is expected?
e) If the amount of CaSO4 measured in the experiment was 3.99 g, what is the percent error?
f) What is the percent yield?

48. 50A. c)SO42- + Ca2+  CaSO4(s) d)4.3 g CaSO4 e)-7.2% f)93%
a) CuSO4 (aq) + Ca(NO3)2 (aq)CaSO4 (s) + Cu(NO3)2 (aq)
b)Cu2+ +SO42-+ Ca2+ + 2NO3-  CaSO4 (s) +Cu2+ + 2NO3-
c)SO42- + Ca2+  CaSO4(s)
d)4.3 g CaSO4
e)-7.2%
f)93%