Download Chapter 13…States of Matter

Chapter 11...The Mole
Define:
1. Mole: SI unit for the measurement of amount of a substance
2. Molar mass: mass of element or compound per one mole of substance
3. Empirical formula: smallest ratio of elements in a compound
4. Molecular formula: actual ratio of elements in a molecular compound
Fill in the following mole map:
4.
6.
Mass
Moles
Atoms/Molecules
7.
5.
8.
4.
5.
6.
7.
8.
9.
multiply by molar mass
divide by molar mass
divide by 6.02 x 1023
multiply by 6.02 x 1023
multiply by 22.4L
divide by 22.4L
9.
Liters
Calculate the following:
5. How many moles of aluminum are in 162 grams of aluminum?
6 moles
6. What is the molar mass of CaMg(CO3)2?
184g
7. How many molecules of NaOH are in 112 liters of NaOH?
3.01x1024 molecules
8. How many grams of N2 are present in 67.2 liters of N2?
84g
9. How many molecules of Ca are present in 320 grams of Ca?
4.81x1024 molecules
10. What is the percent composition of carbon in C6H12O6?
40%
11. How many moles of oxygen are in three moles of Mg(NO3)2?
18 moles
Determine the empirical formula for the following:
12. 22.1% aluminum, 25.4% phosphorus, 52.5% oxygen:
AlPO4
Determine the molecular formula for the following:
13. What is the molecular formula of a compound that has 40.00% carbon, 6.67% hydrogen and 53.33%
oxygen, and has a molar mass of 150 amu’s?
C5H10O5
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Chapter 12…Stoichiometry
Use the following reaction to answer the questions that follow:
Fe(s) +
O2 (g)
Fe2O3 (s)
1. How many grams of Fe2O3 will be produced from 213 grams of Fe?
304.29g
2. How many liters of O2 are required to produce 96 grams of Fe2O3?
20.16L
Fe(s) +
O2 (g)
Fe2O3 (s)
3. How many grams of iron are required to react with 82 liters of O2?
273.33g
4. In the following reaction: H2 (g) +
produce 16.8 liters of NH3?
25.2L
N2(g)
NH3 (g) How many liters of H2 are required to
5. A sample of 23.4 grams of sodium hydroxide is mixed with 18.7 grams of lithium sulfide, producing lithium
hydroxide and sodium sulfide. What is the limiting reagent? How many grams of lithium hydroxide are
produced?
14.01 g
6.
LiOH + KCl  LiCl + KOH
a) I began this reaction with 20 grams of lithium hydroxide. What is my theoretical yield of lithium
chloride?
35.4 g
b) I actually produced 6 grams of lithium chloride. What is my percent yield?
16.9%
Chapter 13…States of Matter
Define:
1. Vapor Pressure: pressure exerted by a vapor over a liquid
2. Freezing Point: matter becomes a solid; liquid to solid
3. Melting Point: matter become a liquid; solid to liquid
4. Boiling Point: matter becomes a gas; liquid to gas
5. Condensation Point: matter becomes a liquid; gas to liquid
6. Sublimation: solid to gas
8. Triple Point: all three phases exist at same temperature and pressure
5. Tyndall Effect: scattering of light by colloid particles
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Pressure (mm Hg)
Use the Phase Diagram Below to Answer the Following Questions:
55,327
Liquid
Solid
3885
Gas
760
Vapor
-57
-78
31
Temperature (C)
1. What is the triple point of the substance depicted in the phase diagram above? -57
2. What is the melting point? ___-57________ freezing point? ______-57_________
3. What phase is this substance in at –67C and 3900 mm Hg of pressure? solid
Hydrogen bond
dispersion forces
dipole-dipole
Chapter 14…Kinetic Theory
Define:
1. Barometer: an instrument that is used to measure atmospheric pressure
2. Temperature: a measure of the average kinetic energy of the particles in a sample of matter
3. Gas Pressure: # of collisions on the wall of the container
4. Kinetic Energy: energy in motion
5. Heat: a form of energy that flows from a warmer object to a cooler object
6. Absolute Zero: 0 K; where all motion stops
Short Answer:
7. Relationship between volume and # of collisions: less volume more collisions
8. Relationship between temperature and # of collisions: more temperature more collisions
9. Relationship between pressure and # of collisions: more pressure more collisions
10. Relationship between volume and pressure: less volume more pressure
11. Relationship between temperature and volume: more temperature more volume
12. Heat always flows from __high______ to __low______.
Temperature Conversions:
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12. 334 K = __61______ °C
Pressure Conversions:
14. 0.96 atm = __97.25__ kPa
13. 25°C = __298_____ K
15. 324 kPa = __2430___ mmHg
16. 956 mmHg = ___126___ atm
Chapter 14 cont...Gas Laws
Write the equations for the following and indicate what is held constant:
1. Boyle’s Law: P1V1=P2V2
2. Charles’ Law: V1 = V2
T1 T2
3. Gay-Lussac’s Law: P1 = P2
T1 T2
4. Combined Gas Law: P1V1 = P2V2
T1
T2
5. Ideal Gas Law: PV=nRT
6. Density: D= mass/volume
7. Indicate the relationship between the variables in each of the equations above.
As temp increases, pressure increases
As pressure increases, volume decrease
As temperature increases, volume increases
As pressure increases, density increases
As volume increases, density decreases
8. Indicate what each variable stands for in each of the equations above.
P=Pressure; V=Volume; T=Temperature; n=moles; m=mass; R=Gas Constant (8.31Kpa/molK or
.082atm/molK)
9. Temperature is in what unit for gas laws?
Kelvin
10. STP stands for: Standard temperature and pressure
11. Standard temperature = _0oC or 273 K___
12. Standard pressure = ___1____ atm, ___101.3____ kPa or __760____mmHg
13. What is the relationship between: volume and density: inverse
temperature and density: inverse
pressure and density: direct
Calculate the following:
14. A sample of argon gas is cooled and its volume went from 830 mL to 520 mL. If its final temperature was
35°C, what was its original temperature?
491.62K
15. A sealed container of nitrogen gas has a pressure of 780 kPa and a temperature of 18°C. The can of gas is
left in a hot car and the temperature of the gas increases to 65°C. What is the new pressure in the can?
905.98kPa
16. A sample of fluorine gas exerts a pressure of 900 mmHg. When the pressure is changed to 1.75 atm, it
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occupies 500 mL. What was the original volume of the gas in liters?
741.53mL
17. How many moles of O2 gas will occupy a volume of 435 mL at 267 kPa and 25°C? How many molecules
of O2?
.05 moles O2
3.01x1022 molecules
18. A 300 mL sample of CO2 is submitted to conditions of 5.2 atm and 33°C. What would the volume of this
gas be at STP?
1391.7mL
Chapter 15…Solutions
Define:
1. Solution: homogenous mixture composed of a solute and solvent
2. Solute: gets dissolved
3. Solvent: does the dissolving
4. Miscible: liquids that are capable of dissolving each other
5. Dilute: when a solution has a low concentration of solute
6. Concentrated: when a solution has a high concentration of solute
7. Saturated solution: cannot hold any more of a given solute at a given temperature
8. Unsaturated solution: can hold more of a given solute at a given temperature
9. Supersaturated solution: can hold more of a given solute than it should theoretically hold (requires heating and stirring)
10. Precipitate: an insoluble solid that will separate from a solution during a reaction (double displacement).
Use the solubility graph below to answer questions 11-14:
11.
12.
13.
14.
15.
16.
How many grams of NaNO3 will dissolve at 30C? ~ 90oC
Which substance is least soluble at 10C? KNO3
Which two substances have the same solubility at 72C? KNO3 & NaNO3
80 grams of KBr placed in 60C creates a (saturated/unsaturated/supersaturated) solution.
When Be(NO3)2 (aq) and NaOH (aq) are mixed together, the resulting precipitate is:
Placing a solute in a solvent will (increase/decrease) the boiling point and (increase/decrease) the freezing
point of a solution.
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17. (Ionic/molecular) compounds cause greater change of boiling/freezing points in solutions.
18. Aluminum chloride (AlCl3) will dissociate into __4___(how many) ions and will cause a (greater/smaller)
change boiling point than MgCl2.
19. In the solvation of solids, solubility rate increases with (high/low) temperatures and (high /low) pressure.
Pressure does not affect solids,only gases
20. In the solvation of gases, solubility rate increases with (high/low) temperatures and (high /low) pressure.
21. The general rule for solvation is “___like_____ dissolves ___like_____”.
For each of the following reactions…
a) Finish and balance the chemical reaction
b) Use a solubility chart to predict what precipitate will form in the reaction
c) Write the complete ionic equation
d) Write the net ionic equation
22. Al(ClO4)3 (aq) + K2CO3 (aq)
a)
b) Al2(CO3)3
c)
d)
23. AgNO3 (aq) + H2S (aq)
a)
b) Ag2S
c)
d)
Compound
Solubility
Salts of group 1A and
ammonia (NH3)
Soluble (aq)
Some lithium compounds
Ethanoates, nitrates,
chlorates, and
perchlorates
Soluble (aq)
Few Exceptions
Sulfates (SO4)
Soluble (aq)
Compounds of Pb, Ag, Hg, Ba, Sr, and Ca
Chlorides, Bromides,
and Iodides
Soluble (aq)
Compounds of Ag and some of Hg and Pb
Insoluble (s)
Group 1A compounds. Compounds of Ba, Sr,
and Ca (slightly)
Insoluble (s)
Group 1A and ammonia (NH3) compounds
Sulfides and
Hydroxides
Carbonates,
Phosphates, and
Sulfites
Exceptions
Chapter 16…Thermodynamics
Define and give the symbols for the following:
1. Entropy: the measure of the amount of disorder in a system.
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2. Heat of reaction (enthalpy): the internal heat content of a substance; H – change in heat
4. Heat of formation: the heat released or absorbed (enthalpy change) during the formation of a pure
substance and usually denoted by ΔHf.
6. In general, spontaneous reactions favor (endothermic/exothermic) reactions and a(n) (decrease/increase) in
entropy.
8. Rank the four states of matter (2.gas, 3.liquid, 1.plasma, 4.solid) in order of decreasing entropy:
9. In the reaction: NH3 (g) + HCl (g)
NH4Cl (s) entropy (decreases/increases).
10. Standard conditions in thermodynamics are: ___25_ C, ___298___ K, ___1___ atm, __101.3____ kPa and
_760__ mmHg.
Label the Following Lines on the Heating Curve as: Boiling Pt., Melting Pt., Liquid, Vapor or Gas:
Temp.
Gas/ Vapor
Liquid
Solid
Boiling point/ Condensation Point
Melting point/ Freezing point
Time
4. Calculate the amount of energy required to heat a 150 g chunk of aluminum from 20C to 40C. (Cp of
aluminum = 0.220 cal/gC)
H=mCpT (150g)(.22)(20) = 660 cal
Chapters 17& 18…Reaction Rates & Equilibrium
Define:
1. Equilibrium: the reaction occurs simultaneously in both directions.
2. Activated complex: Temporary unstable cmpd that is the transition between reactant and product; May form products
or may break apart to re-form products; Also called “transition state”
3. Activation energy: the minimum amount of energy particles need to react with each other.
4. Catalyst: something that speeds up a reaction without being used up.
5. Inhibitor: substance that acts against the effects of a catalyst.
7. The reaction rate expression for an irreversible reaction is: Rate = __k [A]m [B]n
8. The equilibrium constant expression for a reversible reaction is Keq = __[C]c [D]d
[A]a [B]b
10. List the five factors that can affect the rate of reaction:
a. nature of reactants
b. temperature
d. particle size/surface area
e. catalyst
c. concentration
f. inhibitor
Use the equation below to answer the following questions:
H2 (g) + Br2 (g)
2 HBr (g) + Heat
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12. What is Keq if the concentrations of H2 and Br2 are each 2.0 M, and the concentrations of HBr is 3.0 M?
13.
14.
15.
16.
17.
18.
Do more products or reactants exist at equilibrium? Rate is the same, concentration doesn’t matter
If more H2 is added to this reaction, equilibrium will shift (left/right/no shift).
If Br2 is removed from this reaction, equilibrium will shift (left/right/no shift).
If pressure is increased, equilibrium will shift (left/right/no shift).
If temperature is decreased, equilibrium will shift (left/right/no shift).
If HBr is increased, equilibrium will shift (left/right/no shift).
20. A catalyst (does/does not) become part of a chemical reaction.
55)
According to the Common Ion effect, what will happen at equilibrium to the equation: BaSO4(s)
+ Na2SO4(aq) →
The rate will increase to form more products
56)
List the 3 reasons some rxns are not reversible.
Product is a gas
Product is a precipitate
Product can’t ionize
Chapter 19…Acids, Bases and Salts
1. Acid: Compound that produces H+ when dissolved
2. Base: Compound that produces OH- when dissolved
3. Salt: any ionic compound that is not an acid or a base
4. Amphoteric: a substance that can act as either an acid or a base
5. Polyprotic acid: an acid with more than 1 H+
6. Binary acid: An acid that has 2 elements
7. Ternary acid: An acid that has more than 2 elements
8. Acidic anhydride:
9. Basic anhydride:
10. Electrolyte: A substance that can conduct electricity when dissolved in water
Fill in the following table of acid-base theories:
Acids
Bases
Arrhenius Theory
Gives off H+
Gives off OH-
Bronsted-Lowry Theory
Proton ion donor
Proton ion acceptor
Lewis Theory
Electron acceptor
Electron donor
Name the following:
15. HC2H3O2: acetic acid
16. HF: hydrofluoric acid
17. H2SO3: sulfurous acid
Write the formulas for the following:
18. hydrosulfuric acid: H2S
19. perchloric acid: HClO4
20. nitrous acid: HNO2
21. Label the compounds in the following reaction as conjugate acid (CA), conjugate base (CB), acid and base:
OH (aq) + H2CO3 (aq)
Base
Acid
HCO3 (aq) + H2O (l)
CB
CA
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23. Water ionizes into ____H+______ and ____OH-______.
24. The value for Kw is always equal to ___1 x 10-14____________.
25. If the concentration OH- is 6.3 x 10-3, what is the concentration of H3O+ in water?
1.58 x 10-12
26. Is the above solution acidic or basic?
Basic
27. What are the products of a neutralization reaction?
Water and salt
Use the following concentration to answer questions 26-29:
**[H+] concentration of 3.62 x 10-12
28. pH = 11.4
29. pOH = 2.6
30. [OH-] concentration = 2.5 x 10-3
31. Is this solution acidic or basic? Basic
Chapter 22...Organic- Hydrocarbons
Define:
1. Hydrocarbon: Covalent compound made of hydrogen and carbon
2. Alkane: single bonds only; saturated
3. Alkene: at least 1 double bond; unsaturated
4. Alkyne: at least 1 triple bond; unsaturated
5. Functional group: substituent off the parent chain
Identify the following prefixes:
12. one carbon: meth
13. two carbons: eth
14. three carbons: prop
15. four carbons: but
16. five carbons: pent
17. six carbons: hex
18. seven carbons: hept
Identify the following as an
alkane, alkene, or alkyne:
22. C3H6: alkene
23. C6H10: alkyne
24. C2H6: alkane
19. eight carbons: oct
20. nine carbons: non
21. ten carbons: dec
Name the following:
25. Alkyne with 9 carbons: nonyne
26. H H H H
| | | |
C=C-C-C-H
|
| |
H H H
Identify the following
functional groups:
29. R–OH: alcohol
30. R-O-R: ether
31. R-NH2 amine
27. C8H18: octane
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1-butene
O
||
-C-H
31. R-CHO or
aldehyde
32. R-CO-R or
ketone
28.
cyclooctane
O
||
C-C-C
33. R-COOH or
Carboxylic acid
34. R-COO-R or
ester
O
||
-C-OH
O
||
C-C-O-C
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