Download Review Chem Final Review Problems for Final Exam Neil

Name____________________________________Period_____
Review Problems 2013
NOTE: These problems in no way represent all of the problems or concept on the test! This is just a place for
you to start studying!
Naming Rules:
a)
b)
c)
d)
e)
f)
How are compounds with a metal named differently than those with non-metals?
What are the differences between chemicals that end in –IDE, -ITE, and –ATE?
What are the polyatomics you are responsible for memorizing?
How do you name a chemical with a transition metal in it? (You need to know rules for Copper and Iron)
When do you use prefixs like mono, di, and tri? What is the ending for the first element? For the second
element?
How are acids without oxygen named? With oxygen?
Naming Practice:
Give the name for the following formulas:
1) Al(NO3)3
2) BeO
3) NH4Cl
4) Fe2O3
5) P2O5
6) Li2S
7) H2SO4
8) H2S
9) (NH4)3P
10) LiCl
11) Cu2S
12) CaF2
13) H3PO4
14) NaClO3
15) N2O
16) Ba(OH)2
17) HCl
18) Mg3N2
Give the formula for the following names:
1) Magnesium Carbonate
2) Potassium Phosphide
3) Tin (II) Oxide
4) Dinitrogen Tetrafluoride
5) Cupric Chloride
6) Calcium Hydroxide
7) Ferric Nitride
8) Sulfur Dioxide
9) Sodium Sulfite
10) Hydronitric Acid
11) Magnesium Nitrate
12) Barium Iodide
13) Nitrous Acid
14) Potassium Acetate
15) Calcium Chloride
Numbers Rules:
a) Captive Zeros are _______ significant. Leading Zeros are _______ significant. Trailing Zeros are _______
significant, based on the location of the _____________.
b) Adding/Subtracting with sig figs the answer is rounded to the least number of ___________.
c) Multiplying/Dividing with sig figs the answer is rounded to the least number of __________.
Numbers Practice:
1) How many significant digits?
a) 30,100
b) 102.010
c) 0.00270
d) 4,010.
e) 0.05060
f) 0.90
g) 10050
h) 0.760
i) 101.00
j) 200.034
2) Answer the following problems using significant figures
a) 2.65 – 1.9 =
b) 150.0/2.7 =
c) 0.067 + 1.01 + 2.5 =
d) 8.7 + 15.43 + 19 =
e) 1.50/2 =
f) 8.08 x 5.3200 =
3) Write the following number in scientific notation:
a) 0.000398
b) 53,700,000
c) 43,001,000,000
d) 0.010
Rules for Moles:
a) How many atoms are in one mole of an element?
b) How many molecules are in one mole of an element?
c) How do you determine the number of grams in an atom or molecule?
Moles Practice:
1) How many atoms are in 5.2 g of sulfur?
2) How many moles are in 4.2 x 1025 molecules of HBr?
3) How many molecules are in 3 moles of C6H12O6? How many grams?
4) How many moles are in 8.32 x 1020 atoms of Calcium? How many grams?
5) How many atoms of Oxygen are in 3 molecules of C6H12O6?
6) How many moles are in 7.2 grams of Al2(SO4)3? How many molecules? How many atoms of Sulfur?
Important Vocabulary about Atoms: Write definitions to the following in your own words:
1) Atomic Number
2) Atomic Mass
3) Ion
4) Isotope
5) Cation
6) Anion
7) Proton
8) Nuetron
9) Electron
10) Atom
11) Molecule
Reactions in Solutions and Stoichiometry:
Complete and balance the following equations by predicting the products of each chemical reaction.
1) Single Replacement:
Na(s) + CuSO4(aq) 
3) Decomposition:
K2O(s) 
2) Double Replacement :
CuCl2(aq) +
Na3PO4(aq) 
4) Single Replacement:
Al(s) + FeCl2 
5) Be able to define/give examples of the following types of reactions:
a. Precipitation
c. Oxidation-Reduction
e. Double-Replacement
g. Decomposition
b. Acid-Base
d. Single Replacement
f. Synthesis/Combination
h. Combustion
6) Aqueous solutions of copper (II) sulfate and sodium sulfide are mixed.
a. Write a molecular equation for this reaction. Include physical states for all chemicals
b. Write a complete ionic equation for this reaction. Include physical states.
c. Write a net ionic equation for this reaction. Include physical states.
7) Aluminum reacts with iron (III) oxide to produce aluminum oxide and iron. Write and balance the equation.
8) Nitrogen gas reacts with hydrogen gas to produce ammonia (NH3). Write the correct equation and balance it.
9) Water is formed when oxygen(O2) and hydrogen (H2) gas combine. How many moles of water are produced when 5
moles of oxygen react with more than enough hydrogen?
10) Use equation from 9. If 19 grams of hydrogen gas reacts with more than enough oxygen, how many grams of water
will be formed?
11) 2HCl + MgO  MgCl2 + H2O
a.
b.
c.
d.
How many grams of water is formed when 28 grams of HCl react?
How many grams of MgO are needed to make 10 grams of water?
If the mass of water produced is 500 grams, how many grams of HCl were reacted?
What mass of water is produced when 11.2 grams of MgO react?
12) C25H52 + 38O2  25CO2 + 26H20 If 100 grams of oxygen is burned, how many grams of carbon dioxide are produced?
13) 2KClO3  2KCl + 3O2 How many grams of KCl are produced when 48.2 grams of potassium chlorate react?
14) 2H2 + O2  2H2O
a. What is the limiting reactant if 4g of hydrogen reacts with 16g of oxygen.
b. What mass of water is produced when 9g of hydrogen reacts with 50g of oxygen?
c. What mass of water would be produced if 2g of hydrogen reacts with 11g of oxygen?
Gas Laws:
1. A 3.00-g sample of KClO3 is decomposed and the oxygen at 24.0°C and 0.982 atm is collected. What volume of
oxygen gas will be collected?
2KClO3(s)  2KCl(s) + 3O2(g)
2. The volume of a sample of gas is 650. mL at STP. What volume will the sample occupy at 0.0°C and 950 mmHg?
3. A helium balloon has a volume of 2.30 L at 23.5°C and a pressure of 1.00 atm at sea level. The balloon is released and
floats upward. At a certain height the atmospheric pressure is 0.810 atm and the temperature is 12.0°C. Calculate the
volume of the balloon.
4.A sample of carbon monoxide is collected at 55°C and 0.892 atm. What will its new pressure be at 20°C?
5.A sample of helium gas occupies 2.65 L at 1.20 atm. What pressure would this sample of gas exert in a 1.50-L container
at the same temperature?
6.What volume of chlorine gas, measured at STP, is necessary for the complete reaction of 4.81 g of sodium metal?
2 Na (s) + Cl2 (g)  2 NaCl (s)
7. A mole of a gas is at 1.7 atm and 37°C and the pressure is changed to 813 mmHg, calculate the new temperature.
8. You transfer a sample of a gas at 17°C from a volume of 5.67 L and 1.10 atm to a container at 37°C that has a pressure
of 1.10 atm. What is the new volume of the gas?
9. What volume is occupied by 19.6 g of methane, CH4, at 27°C and 1.59 atm?
10. What is the mathematical relationship (inverse vs. direct) between all of the different combinations of variables in
gas laws – volume, pressure, temperature, and moles? (For example between volume and pressure or volume and
temperature.)
Equilibrium
1. Given the following equation: H2(g) + I2(g)   2 HI(g)
a. Write the equilibrium expression.
b. The Keq for this equation is 2.4 x 10-5. Determine Q given the following information: 4.5647 x 10-3 M of H2,
7.378 x 10-4M of I2 and 1.3544 x 10-2 M of HI.
c. Which direction will the reaction have to proceed in order to reach equilibrium?
2. Consider the following equilibrium system:
C(s) + CO2(g) + heat  2 CO(g)
If the reaction is at equilibrium, what would be the effect (how would it shift) of:
a) adding CO2(g)
b) adding C(s)
c) adding heat
d) increasing the pressure of the system by decreasing the volume
e) removing CO(g)
3. At equilibrium explain what happens to:
a) The concentration of products and reactants
b) The rate of the forward and reverse reaction.
Acids & Bases
1. What is the difference between the Bronsted-Lowry and Arrhenius definitions of acids and bases?
2. What is the difference between a strong acid/base and a weak acid/base? Are strong acids strong or weak
electrolytes?
3. Strong acids have a _________ (weak or strong) conjugate base. Weak acids have a ____________ (weak or
strong) conjugate base.
4. How is the pH scale determined? What is the relationship between pH and pOH?
5. How do you identify a substance as an acid, base, or neutral from its pH? From its pOH? From its [H+]? From its
[OH-]?
6. For each of the following calculate the indicated quantity, also indicate whether the solution is an acid or base.
a. [OH-]=6.62x10-3M , pH=?
b.
pH =6.325, [OH-] = ?
c.
pH = 9.413, [H+] =?
d. [H+] = 2.45x10-10 M, [OH-] = ?
7. What are indicators and why are they important?
8. What is the equivalence point for each of the following titrations – also be able to explain why the equivalence
point is located at that pH.
HCl & NaOH
b. HC2H3O2 and NaOH
c. HCl and NH3
d. HC2H3O3 and NH3
9. How many grams of Al(OH)3 are needed to react with 295 L of an HNO3 solution that has an initial pH of 2.4? The
molar mass of Al(OH)3 is 78.03 g/mol.
Molecular geometry
1. How do electronegativity differences between two atoms determine the type of bond they will form?
2. Complete the lewis dot structure for the following and then categorize them as: linear, trigonal planar,
tetrahedral, pyramidal or bent
a. NF3
b. BF3
c. CCl4
d. H2S
e. CO2
3. Describe the structure of the atom according to the most recently accepted model.
4. Write the orbital diagrams for the following
a. Co
b. Se
c. Ar
5. How many “s” electrons does Oxygen have?
Solutions
1. Explain the difference between these terms: Saturated, Unsaturated, Dilute, and Concentrated
2. Explain the phrase “like dissolve like,” especially in terms of the polarity of compounds.
3. What 3 things can you do to increase the solubility of a “salt?”
4.
Given the following equation: Mg(NO3)2(aq) + 2 KCl(aq)  MgCl2(s) + KNO3(aq)
What mass of MgCl2 is produced when 258 mL of a 0.25 M solution of KCl is reacted with 325 mL of a 0.145 M
solution of Mg(NO3)2?
5.
You want to make 250 mL of a 0.15 M HCl from a 0.5 M solution of HCl. What volume of the 0.5 M HCl do you
need to start with?