Download review sheet

REVIEW SHEET SEMESTER TWO
PREAP CHEMISTRY
Bonding
1.
2.
3.
4.
.In ionic bonding electrons are ___________________.
In covalent bonding electrons are ________________.
Why do atoms bond?________________________________________________
Using electronegativities determine what type of bonding, molecular shape and molecular polarity for
the following molecules.
Compound
Type of bond
Molecular shape
Molecular polarity
NaCl
CH4
H2O
NH3
AlCl3
Kinetic Theory and States of Matter
1. Name the four states of matter.
2. Name the three intermolecular forces.
Which is the strongest ?_________________ Which is the weakest?________________
3. What are the four elements necessary for H-bonding?
4. Arrange the following molecules in order of decreasing boiling point.
CH4
H2S
N2
NH3
5. Why does ice float?
6. At what temperature is water most dense? Why?
7. What is rate of flow?
8. What factors affect rate of flow?
9. What substance will decrease the surface tension of water?______________________
10. A highly volatile substance will have _______ intermolecular forces.
11. What type of intermolecular force will be the strongest in polar molecules?________________Nonpolar
molecules?_______________
12.Adding heat to a solid will ____________ increase/decrease the kinetic energy and cause the intermolecular
forces to ________________.
13. Of the three phases of matter, ______________ has the strongest attractive forces, and
_________________ has the weakest.
14. Where will water boil at a lower temperature 250 feet or 10,000 feet above sea level?
_____________________________________________ Why?_____________________
15. Explain on the molecular level how solids, liquids and gases compare. Which has the most kinetic energy?
16. Label the heating curve of water wit liquid, solid, gas, boiling point, freezing point, melting point,
condensation point. Show in which direction energy increases.
17. How much energy would it take to melt 10 moles of ice if the molar heat of fusion for water is 6.009
KJ/mole?
18. q=m(ΔT) Cp How much heat does a 100g cube of ice need to go from -200Cto 00CC? The specific heat of
ice is 2.06 J/g 0C.
19. Draw a phase diagram for CO2 and label the normal melting point, normal boiling point, triple point,
critical temp and critical pressure.
The Gas Laws
1. If pressure decreases, temperature _____________, and volume ___________________.
2. The number of collisions of gas molecules is measured by ________________. List three ways to increase
the number of collisions.
____________________________________
___________________________________
____________________________________
3. Gases have _______________kinetic energy, __________________distances between molecules and
therefore ____________________ compressed.
4. Standard temperature is ___________________0C and ____________________K
5. Convert.
a. 160C = _____________________K
b. 16K = _____________________0C
6.
a.
b.
c.
d.
e.
Convert
0.95 atm = ________________________mmHg
790 mmHg = ______________________kPa
125 kPa= _________________________mmHg
100 mmHg = ______________________ Torr
95 kPa = __________________________atm
7. If gas A travels slower than gas B at 350C, then what can be said about the mass of A and B?
8.
LAW
Charles’
CONSTANT
FORMULA
Bolye’s
Gay-Lussac’s
Ideal
N/A
Combined Gas
N/A
9. At 1120C, the volume of a gas is 2.3L. Find the volume if the temperature is changed to 650K and the
pressure remains the same.
10. A gas collected when the pressure is 80 kPa has a volume of 175 mL. What volume will the gas occupy at
standard pressure?
11. At 327K, a gas exerts a pressure of 0.7 atm. If the volume remains constant, what will happen to the
pressure if the temperature is raised to 876K.
12. What is the volume, in liters, of 0.250 mol of oxygen gas at 20.00C and 0.974 atm pressure?
Solutions
1. What are the three factors affecting rate of solubility?
2. Define
Saturated
Unsaturated
Supersaturated
3. What is meant by dissociation?
4. In a solution, the ___________________ is dissolved in the ____________________.
5. When two substances will not dissolve in one another the are said to be _____________. When they will the
are _________________________.
6. Calculate the molality (m) prepared by dissolving 15 g of NaCl in 500 g of water.
7. A 250 mL solution has a molarity (M) of 2.5 M. How many moles of solute are contained in this solution?
8. A solution contains 10.0 g of CuSO4 in 100mL of solution. What is the percent (m/v) of the solution?
9. 10 mL of a 3M solution of KCl is diluted with 350mL of water. What is the molarity of the diluted
solution?
Reaction Rates and Equilibrium
1. List the parts to the collision theory.
2. In the following reaction mechanism, identify the intermediates, and give the net equation.
Step 1 2NO(g)  N2O2(g)
Step 2 N2O2(g) + O2(g)  2NO2(g)
3. What is a catalyst?___________________________________________
4. What kind of catalyst is in your car?_____________________________
5. What is an inhibitor?_________________________________________
6. List the five factors that influence reaction rates.
___________________________________
__________________________________
____________________________________
___________________________________
___________________________________
7. What is the sixth factor that only applies to gases?___________________________
8. In the following one-step reaction give the rate law.
2NO(g) + O2(g)  2NO2(g)
What order of reaction is it overall?
In terms of NO?
In terms of O2?
9. Given the following reaction mechanism, write the rate law.
2NO  N2O2
FAST
N2O2 + O2  2NO2 SLOW
2NO + O2  2NO2
10. What is the rate determining step for the above reaction?
11. Draw and label an exothermic and an endothermic energy diagram. Label them with Ez, Ez’, ΔE, reactants,
products, and activated complex. Make sure and label whether ΔE is positive or negative.
12. Calculate the equilibrium constant Keq for the following reaction.
2A <=> 2C + D
At equilibrium, the concentration of A is 6.59 X 10-4 M, the concentration of C is 4.06 X 10-3 M, and the
concentration of D is 2.03 X 10-3 M.
13. Explain how a system at equilibrium responds to a stress and list factors that can be stresses on an
equilibrium system.
14. Explain how decreasing volume of the reaction vessel affects each equilibrium.
a. 2SO2(g) + O2(g) <-- > 2SO3(g)
b. H2(g) + Cl2(g) < -- > 2HCl(g)
15. Decide whether higher or lower temperatures will produce more CH3CH0 in the following equilibrium.
C2H2(g) + H2O(g) < --> CH3CHO(g)
∆H = -151 kJ
16. Make sure you can calculate rate law from data given!
Thermochemistry
1. Calculate the Gibb’s Free Energy change for the following reaction at 25oC.
Ca(s) + 2H2O(l)  Ca(OH)2(s) + H2(g).
Given Hrxn= -411.6 kJ/mol; Srxn=.0318 kJ/mol (K)
2. Give the formula for Hreaction and state Hess’s Law.
3. Use Hess’s law and the changes in enthalpy for the following two generic reactions to calculate ∆H for the
reaction 2A + B2C3  2B + A2C3.
2A + 3/2C2  A2C3
∆H = -1874 kJ
2B + 3/2 C2  B2C3
∆H = =285 kJ
4. Use standard enthalpies of formation from the table to calculate ∆Hrxn for the following reaction.
P4O6 (s) + 2O2(g)  P4O10(s)
5. Is the following reaction endothermic or exothermic. How much energy is gained or lost if 2 moles of iron
reacts completely with the oxygen?
4Fe(s) + 3O2(g)  2Fe2O3(s) + 1625kJ
5. How can we predict if a reaction is spontaneous? Complete the table.
∆H
∆S
∆G
Negative (-) value
Positive (+) value
Negative (-) value
Negative (-) value
Positive (+) value
Positive (+) value
Positive (+) value
Negative (-) value
Acids and Bases
1. List 4 properties of acids and 4 properties of bases.
2. Define an Arrhenius Acid and an Arrhenius Base
3. Define a Bronsted-Lowry acid and base.
4. Define a Lewis acid and base.
5. The pH of an acid is ____7, a base is _____7 and a neutral solution is _______7.
6. Give the formula to find pH.
7. How do you find pH if you are given [OH-] ?
8. What are the products of a neutralization reaction?
9. Label the acid, base, conjugate acid, and conjugate base for the following equation.
NH3 + HNO3  NH4+ + NO3-
10. Find the pH of the following and label as acid, base(alkaline), or neutral
[H+]= 1.0 x 10-5
[H3O+] = 2.0 x 10-13
[OH-] = 3.0 x 10-4
11. Given pH = 4.3, find pOH, [OH], [H3O+1]
12. Write formulas for and give names for the following acids.
H2SO4
HNO3
HF
HI
Phosphoric acid
Sulfurous acid
13. Write formulas for and give names for the following bases.
NaOH
Ca(OH)2
NH3
Barium hydroxide
Magnesium hydroxide
14. If 20.00 mL of a 0.01 M solution of HCl is titrated with NaOH, 15.00 mL of NaOH is used at the endpoint.
What is the molarity of the base?
15. What is the Ka of an acid that has a [H+] of 2.5 x 10-3M and the concentration of athe acid is .2M?
16. If the concentration of [Ag+1] is 2.53 x 10-4 M, what is the Kspof Ag2CO3?
REDOX
1. Give the oxidation numbers of all the elements in the following compounds.
KCN
MgSO4
PO4-3
CaCO2
2. Identify the following as a redox or nonredox
H2  2H
2Na + Cl2  2 NaCl
HCl + NaBr  HBr + NaCl
3. The oxidation number of a free element is always ____________________.
4.The most active reducing agent among the elements is ____________________.
5. For each of the following equation s determine whether oxidation or reduction is happening.
K  K+ + eS + 2e-  S-2
6. What is the difference between a Voltaic cell and a electrolytic cell.?
7. Calculate Eocell for the following.
a. Na+/Na and Ni2+/Ni
b. Cu2+/Cu and Ag+/Ag
8. Identify the compound being oxidized, reduced, the oxidizing agent, and the reducing agent. Write the half
reactions and balance.
Sb + H2SO4  Sb2(SO4)3 + SO2 + H2O
9. Draw and label a voltaic Cell with the following cell notation; Zn/Zn+2//Cu+2/Cu
10. Draw and label an electrolytic cell that shows silver electroplating of a nickel ring.