### semester two final review key units 7 and 8

```SEMESTER TWO FINAL REVIEW SHEET
Below are the types of questions you can expect to see on your final for Fundamentals. The topics are
arranged in the order that we learned the concepts. If there are problems you do not understand, or need
additional practice on, please revisit that unit’s review sheet from my website.
In addition to a periodic table, scientific calculator and ion chart, you may also use this HANDWRITTEN
review sheet. You may choose to answer directly on this sheet or on a separate sheet of paper and attach it.
UNIT 7: KINETIC MOLECULAR THEORY AND GAS LAWS
KINETIC THEORY, STATES OF MATTER, TEMP CONVERSION
1. Definitions to Know: Solid, Liquid, Gas, Phase change, temperature, heat, pressure, and absolute zero.
Solid: one of the four fundamental states of matter (the others being liquid, gas, and plasma). It is
characterized by structural rigidity and resistance to changes of shape or volume
Liquid: nearly incompressible fluid that conforms to the shape of its container but retains a (nearly)
constant volume independent of pressure
Gas: One of four main states of matter, composed of molecules in constant random motion. Unlike a
solid, a gas has no fixed shape and will take on the shape of the space available.
Phase change: A transition between solid, liquid, and gaseous phases which involves energy
Temperature: Measure of the average kinetic energy of the particles in a substance
Heat: A form of thermal energy that is present in all matter
Pressure: force per unit area
Absolute zero: the temperature where all molecular motion stops
2. What does STP stand for and what are its numerical values? Standard Temperature and Pressure, 1 atm
and 273 K
3. What are the three assumptions of the Kinetic Theory?
(A) All matter consists of small particles
(B) These particles are in constant motion
(C) Collisions between particles are perfectly elastic
1. What are the two temperature scales? Kelvin and Celsius (Farenheit also, but we don’t use it)
5. Convert the following temperatures from Celsius to Kelvin and Kelvin to Celsius:
a. 25°C 298 K
d. 130 K -143°C
b. 333°C 606 K
e. 40 K -233°C
c. -75°C 195 K
f. 289 K 16°C
PHASE CHANGES AND PHASE DIAGRAMS
1. Be able to label the phases/phase changes on a diagram. As practice, fill in the chart below:
PHASE CHANGE
FROM______ TO _______
ENERGY RELEASED/GAINED
EXO/ENDOTHERMIC
EX: Melting
From Solid to Liquid
Energy gained
Endothermic
Freezing
From Liquid to Solid
Energy released
Exothermic
Boiling
From Liquid to Gas
Energy gained
Endothermic
Condensation
From Gas to Liquid
Energy released
Exothermic
Sublimation
From Solid to Gas
Energy gained
Endothermic
Deposition
From Gas to Solid
Energy released
Exothermic
1. Know the boiling point and freezing point of water. BP 100°C and FP 0°C
2. If given a heating curve like the one below, be able to identify when melting and freezing occur and
answer the following the questions.
a. What is the melting temperature of the
substance? 0°C
b. What is the boiling temperature of the substance?
G
100°C
c. What is happening to the temperature at 1, 3,
and 5? It is increasing
d.
What is happening at point 2 and point 4? Phase
L
change
e. Label solid, liquid and gas on the curve.
f. On which point of the curve are the particles
S
moving the least? 1
g. Where on the curve is ONLY a liquid present? 3
h. Based on the melting point and boiling point,
could this substance be water? yes
MIXED GAS LAW PROBLEMS
“R” Constants: 0.0821 L•atm/M•K (For atm)
8.315 L•KPa/M•K (For kPa)
62.4 L•mmHg/M•K (For mmHg)
1. Equations to Know: Dalton’s Law, Boyle’s Law, Charles’ Law, Gay-Lussac’s Law, and Ideal Gas Law.
Dalton’s Law: Ptotal = P1 + P2 + P3 …;
Boyle’s Law: P1 × V1 = P2 × V2
Charles’ Law: V1 / T1 = V2 / T2
Gay-Lussac’s Law: P1 / T1 = P2 / T2
Ideal Gas Law: PV = nRT
2. Calculate the temperature of a gas if 6.5 moles of it occupies a volume of 1.5 L at 2.4 atm. 6.47 K
3. There is a mixture of three gases in a closed container. Gas A exerts a pressure of 10 psi, Gas B a
pressure of 5.5 psi, and the total pressure is 34.7 psi. What is the partial pressure of Gas C? 19.5 psi
4. A gas occupies a volume of 300 cm3 at 200 kPa. What volume will it occupy at 101.3 kPa if the
temperature remains constant? 592.3 cm3
5. The pressure of a sample of helium in a 4.00 L container is 5.5 atm. What is the new pressure if the
sample is placed in a 2.00 L container? 11 atm
6. A gas occupies a volume of 4440 cm3 at a temperature of 35°C. Calculate the volume that this gas will
occupy at 74°C if the pressure is held constant (must convert to Kelvin first). 5002.2 cm3
7. A balloon full of air has a volume of 5.0 L at a temperature of 30°C. What is the balloon's volume at
65°C? 5.58 L
8. Calculate the pressure in kPa exerted by a gas if .40 mole occupies 2.5 L at 298 K. 396.4 kPa
9. A gas has a pressure of 2 atm and a temperature of 289 K. If the temperature increases to 400 K, what is
the new pressure if the volume remains constant? 2.77 atm
10. At 315K a gas exerts a pressure on its container of 101.67 kPa. The temperature suddenly drops to 280K,
what pressure does the gas now exert? 90.37 kPa
11. If I have an unknown quantity of gas at a pressure of 2.4 atm, a volume of 27 liters, and a temperature
of 78°C, how many moles of gas do I have? 22.5 moles
UNIT 8: WATER, SOLUTIONS, UNITS OF CONCENTRATION
PROPERTIES OF WATER
1. Define the following terms: polarity, surface tension, vapor pressure, specific heat, and capillary
action.
POLARITY: Polarity is the separation of charges, positive and negative that can describe a bond or an
entire molecule and is caused by differences in electronegativity.
SURFACE TENSION: The tendency for molecules at the surface of a liquid to be pulled inward resulting
in a smooth surface.
VAPOR PRESSURE: The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid
(or solid)
SPECIFIC HEAT: The amount of energy needed to raise the temperature of 1g of substance by 1°C.
CAPILLARY ACTION: The rise of liquids up a narrow tube
2. Draw four water molecules. Label the types of bonds (covalent vs. hydrogen), oxygen atoms,
hydrogen atoms, and respective charges on the atoms.
H = Hydrogen atoms
O = Oxygen atoms
Hydrogen bonds
δ+
δ-
δ+
Covalent bonds
3. Is water polar or nonpolar? Explain. Water is a polar molecule because the oxygen is more
electronegative that the hydrogen
4. Why is water considered the universal solvent? Water is considered the universal solvent because it
has the ability to dissolve many substances.
5. What are the special properties of water and why do they occur? High surface tension, high specific
heat, low vapor pressure, capillarity, less dense in the solid state; these properties are due to the
hydrogen bonds
6. Explain why solid ice is less than liquid water with regard to particle arrangement. Ice molecules align
themselves in a regular lattice rather than more randomly as in the liquid form. It happens that the
lattice arrangement allows water molecules to be more spread out than in a liquid, and, thus, ice is less
dense than water.
7. Why does a substance like sugar dissolve in water, but oil does not?
Water is a polar molecule and sugar is polar (like dissolves like). Oil is nonpolar so it will not dissolve.
SOLUTIONS
1. Define the following terms: solution, solvent, solute, dilute, concentrated, dissociate, immiscible,
solubility, saturated, supersaturated, and unsaturated.
SOLUTION: a homogeneous mixture where one substance is dissolved inside of another.
SOLVENT: the substance that does the dissolving
SOLUTE: the substance that is dissolved
DILUTE: a solution that has excess solvent; the solution has a lower concentration of solute per solvent
CONCENTRATED: a solution that contains more a higher concentration of solute per solvent
DISSOCIATE: when ionic compounds break into their respective ions completely
SOLUBILITY: the measure of the amount of solute that can be dissolved in a given amount of solvent
SATURATED: a solution where the maximum of solute is added to solvent
SUPERSATURATED: a solution where there are more solute particles than are needed to form a saturated
solution
UNSATURATED: a solution where there are less solute particles than are needed to form a saturated
solution
2. Give an example of solid, liquid, and gas solution. Identify the solute and solvent.
Solid: Steel. Solute-carbon, Solvent-iron
Liquid: Soda. Solute-sugar, CO2, etc. Solvent-water
Gas: Air. Solute-O2, CO2, etc. Solvent-N2
3. What are the factors that affect the solubility of a gas? Temperature and pressure
4. A glass of water has 10g of sugar dissolved in it. If more sugar can be added to dissolve in the
water, is the solution unsaturated, saturated, or supersaturated?
COLLIGATIVE PROPERTIES
1. What are colligative properties? The properties that depend on the relative number of solute and solvent
particles in a solution and not the chemical identity.
2. What are four examples of colligative properties and how are they affected (do they
increase/decrease)?
1. Boiling point elevation 2. Freezing point depression 3. Vapor pressure depression 4. Osmotic pressure
3. Give an example of a colligative property in real-life and explain why/how it is used? salting the roads
where it snows so that the freezing point of water is reduced
CALCULATING CONCENTRATION
Common Conversion Factors:
1000 g = 1 kg
1000 mL = 1 L
For water 1 g = 1 mL
For water 1 kg = 1 L
1. Definitions: molarity, molality, concentration, ppm
Molarity: concentration of moles of solute per liter of solution
Molality: concentration of moles of solute per kg of solven
Concentration: amount of solute per given amount of solution
ppm: grams of solute per grams of solutions times a million
2. If 4.7 g of Na2CO3 is dissolved in 1000 mL of water, what is the molarity of the solution? 0.044 M
3. How many moles of NaCl would be needed to make 12.5 L of a 1.50 M solution? 18.75 moles
4. What is the molarity of a bleach solution containing 6 moles of NaOCl per liter of bleach? 6 M
5. What is the concentration in ppm of ethanol in a solution that contains 0.005 g of ethanol dissolved in
11500 kg of water (1 kg = 1g)? 434.8 ppm
6. Calculate the molarity of a 1.60 L solution containing 0.05 moles of dissolved KBr. 0.03125 M
7. If .5g of blood are added to 10.0 kg of water what is the concentration in PPM? (1 kg = 1g) 50 ppm
8. A solution is made up of 120 g NaOH and 440 g water. The total volume is 200.0 mL. Determine the
following:
a. Moles of NaOH 3 moles
d. PPM
2.72 x 105 ppm
b. Moles of H2O 24.4 moles
e.
Molarity 15M
c. Molality 6.81 m
9. What is the molality of a solution that contains
2 moles of HNO3 in 5.500 kg H2O? 0.364 m
10. What mass of water in kg is required to dissolve
3.5 moles of NaCl to prepare a 3.75m solution?
0.93 kg
SOLUBILITY AND SOLUBILITY CURVES
1. What does like dissolves like mean? Polar
dissolves polar and nonpolar dissolves
nonpolar
2. Why doesn’t pressure affect solids and liquids?
Particles are already close together
3. How many grams of NaCl can dissolve in 100 g
of water at 50oC? 38 g
4. For most substances, solubility increases as
temperature increases. What are the
exceptions on the graph below? NH3 and
Ce2(SO4)3
5. What mass of solute will dissolve in 100g of
water at the following temperatures?
a. KNO3 at 30°C __45 g__________
b. NaCl at 30°C ___37 g_________
c. NH4Cl at 70°C ___60 g_________
d. Which of the above three substances is most soluble in water at 15°C. NaCl
6. On a solubility curve, the lines indicate the concentration of a __saturated_____ solution - the maximum
amount of solute that will dissolve at that specific temperature.
7. Values on the graph below a curve represent unsaturated solutions - more solute could be dissolved at
that temperature.
8. Values on the graph above a curve represent supersaturated solutions – no more solute could be
dissolved at that temperature.
9. Using the solubility graph, determine if the following solutions are saturated, unsaturated or
supersaturated. If they are anything but saturated, list two things you can do to make them
saturated (include numbers).
Solution (in 100g
H2O)
10 g of KClO3 at 20oC
o
40 g NaCl at 40 C
o
60 g KNO3 at 40 C
40 g KClO3 at 80oC
Sat, Unsat,
Supersat
+/- how many °C to
make saturated?
+/- how many g to
make saturated?
SUPERSATURATED
Add about 7 °C
Remove 2 grams
SUPERSATURATED
Add about 50°C
Remove 2 grams
SATURATED
N/A
N/A
SATURATED
N/A
N/A
```