Transcript Semester Two Review Packet - Chippewa Falls High School

SEMESTER TWO
REVIEW PACKET
UNIT NINE: STATES OF
MATTER
• Intermolecular Forces – attractive forces between molecules
• Surface Tension – tendency of liquids to minimize surface area
• Viscosity – resistance of a liquid to flow
• Volatile – easily evaporates
• Vapor Pressure – pressure exerted by a liquid’s vapor molecules
• Melting – phase change from solid to liquid
• Freezing - phase change from liquid to solid
• Boiling - phase change from liquid to gas
• Condensing - phase change from gas to liquid
• Sublimation - phase change from solid to gas
• Deposition - phase change from gas to solid
2. Gases would have weak intermolecular forces whereas solids and
liquids would have strong intermolecular forces.
3. The rate of vaporization increases with:
- increase in surface area
- increase in temperature
- decrease in strength of IMF
4. Vapor pressure increases with:
- decreases in strength of IMF
- increase in temperature
5. To make a liquid boil:
- increase the temperature
- decrease the atmospheric pressure
6. To make a gas condense:
- decrease the temperature
- increase the atmospheric pressure
7. How much energy would be needed to vaporize 43.0 g
of water?
8. How much energy would be needed to melt 28.0 g of
ice?
9. What quantity of heat would need to be added to 123.4 g
of water to change its temperature from 52°C to 86°C?
10. Dispersion is a temporary uneven distribution of electrons that increases with
molar mass, symmetry and surface area.
11.
Dipole – Dipole is a permanent distribution of electrons
12.
Hydrogen Bonding occurs when hydrogen is bonded to nitrogen, oxygen,
or fluorine.
13.
Molecular Solids have neutral corners, low melting points, are nonmetal
with nonmetal or a metalloid with a metal or nonmetal.
14.
Covalent Solids have neutral corners, high melting points, are metalloid
with metalloid, or a metalloid with a metal or nonmetal.
15.
Metallic Solids have positive ions and are metals with metals.
16.
Ionic Solids have positive and negative ions at the corners and are metals
with nonmetals.
17.
Why is water unique?
Water has a low molar mass yet is a liquid at room temperature, easily dissolves
polar and ionic compounds, expands upon freezing, and its ice is less dense than
water.
18. What is the percentage of water, by mass, in
FeSO4•7H2O?
UNIT TEN: GAS LAWS
19. Standard pressure is:
1 atm, 101.325 kPa, 760 mmHg, and 14.7 psi
20. Temperature must be in units of Kelvin.
K = 273 + °C
21. What are the relationships between pressure and
volume, pressure and temperature, volume and
temperature, and volume and the number of moles?
- P increases V decreases
- V increases and T increases
- V increases and n increases
- P increases and T increases
22. A gas mixture contains helium, oxygen, and an unknown gas. Calculate
the partial pressure (in atm) of the unknown gas when the partial pressure of
helium is 503 mmHg, oxygen is 214 mmHg, and the total pressure is 1.13
atm.
23. If a balloon occupies 1.49 L at room temperature (25.0°C), what
temperature in °C would you need to increase the volume to 3.07L?
24. If there is a pressure of 536 mmHg in a container and it is heated from
20.0°C to 100.0°C, what is the new pressure?
25. A hand pump with a moveable piston has an applied pressure of 3.08 atm
and a volume of 3.25L. What is the volume if the applied pressure is
decreased to 1.18 atm?
26. If you transfer a gas at room temperature (25.0°C ) from a 5.89 L cylinder to
a 2.41 L container and the pressure increases from 0.653 atm to 1.36 atm,
what is the final temperature in °C?
27. If you blow up a balloon with helium gas and the volume increases from
56.3 mL to 1006 mL and there are initially 0.163 moles, how many moles
need to be added?
28. If 28.4 L of Cl2 react with an unlimited supply of H2 at STP, how many moles
of HCl would be produced?
29. If helium effuses 3.60 times faster than an unknown gas sample, what is
the molar mass of the unknown gas?
UNIT ELEVEN: SOLUTIONS
30. Terms
• Solution – a homogeneous mixture of two or more substances
• Solute – the minority component
• Solvent – the majority component
• Suspension – a mixture from which the particles settle out upon standing
• Colloid – a permanent mixture whose particles are smaller than a suspension
and larger than a solution
• Tyndall Effect – a way to differentiate solutions from colloids and suspensions
based on the scattering of light
• Electrolyte – aqueous solutions containing a solute that dissociates into ions
and conducts electricity
• Solubility – the amount of component that will dissolve in a certain amount of
liquid
• Saturated – holds the maximum amount of solute under the solution conditions
• Unsaturated – holds less than the maximum amount of solute under the
solution conditions
• Supersaturated – holds more than the normal maximum amount of solute
• Concentration – the amount of solute in a solution
• Mass Percent – the number of grams of solute per 100 grams of solution
• Molarity – the number of moles of solute per liter of solution
• Molality – the number of moles of solute per kilogram of solvent
• Colligative Properties – a property that depends on the number of solute
particles and not the type of solute particle (how much you have not what you
have)
• Freezing Point Depression – difference in temperature between the freezing
point of a solution and the freezing point of the pure solvent
• Boiling Point Depression – difference in temperature between the boiling point
of a solution and the boiling point of the pure solvent
31. Which of the following will scatter light?
- colloid and suspension
32. What would happen to the following if more solute is added to
the solution:
- Saturated: nothing will happen
- Unsaturated: solute will dissolve
- Supersaturated: extra solute will disrupt the solution and
precipitate out
33. Solubility depends on:
- identity of solute and solvent
- temperature
- pressure (only gases)
34. Factors that increase the rate of solution are:
- decreasing the particle size
- stirring
- increasing temperature (except for gases)
35. Solubility Curve
a. Pb(NO3)2 at 30°C
- 66 g
b. 12 g of KClO3
- 35°C
c. 70 g of CaCl2 at 16°C
- saturated
36. A soft drink contains 85.2 grams of sucrose (C12H22O11) in a 741 mL
solution. What is the percent mass if the density of the solution is 1.00
g/mL?
37. What is the molarity of a solution that contains 98.5 grams of MgBr2
dissolved to produce a 1.00 L solution?
38. What volume of 12 M HCl do you need to make 500.0 mL of a 3.0 M
solution?
39. Determine the volume in milliliters of 0.225 M KOH solution required to
neutralize 185 mL of 0.125 M HCl. The neutralization reaction is:
KOH(aq) + HCl(aq) → H2O(l) + KCl(aq)
40. Calculate the molality of a solution containing 225 grams of glucose
(C6H12O6) dissolved in 1.25 L of water. (Assume the density of 1.00 g/mL for
water).
41. What is the freezing point depression and the boiling
point elevation of a 2.43 m solution of Al2(SO4)3 in water?
UNIT TWELVE: ACIDS AND
BASES
42. Determine whether the following are acids (A), bases(B), or both
(X).
B
B Tastes bitter
a. ____
pH>7
b. ____
A
A pH<7
c. ____
Tastes sour
d. ____
A
A Blue litmus red
e. ____
Corrosive to skin
f. ____
B
X Electrical Conductor
g. ____
Feels slippery
h. ____
i. ____
B Red litmus blue
43. Name the following formulas:
a. CsOH
- cesium hydroxide
b. H2S
- hydrosulfuric acid
c. HNO3
- nitric acid
d. HNO2
- nitrous acid
44. Write the formulas of the following names:
a. hydrobromic acid
- HBr
b. iron (II) hydroxide
- Fe(OH)2
c. phosphoric acid
- H3PO4
d. chlorous acid
- HClO2
45. Explain how water is amphoteric.
Water can act as an acid and a base. It is able to accept and
donate protons.
46. List the six strong acids: HCl, HBr, HI, HClO4, H2SO4, HNO3
47. List the eight strong bases: LiOH, NaOH, KOH, RbOH, CsOH,
Ca(OH)2, Ba(OH)2, Sr(OH)2
48. The ion product constant for water is Kw = 1.0 x 10-14
49. Calculate the pH, pOH, [H+], and [OH-]. Identify the
type of solution as acidic, basic, or neutral.
50. What are the reactants and products in an acid base reaction?
- reactants: acid and base
- products: water and salt (ionic compound)
51. What is the purpose of titrations?
A titration is an experimental technique used to determine the
concentration of an unknown acid/base by comparing it to an acid/base
solution of known concentration
52. H2SO4 (aq)
+
Ba(OH)2 (aq)
53. 2 Al(OH)3(aq) + 3 H2SO4(aq) →
→ BaSO4(aq) + 2 H2O(l)
Al2(SO4)3 (aq) +
6 H2O(l)
54. What is the molarity of a solution of NaOH if 34.2 mL of the solution
is used to neutralize 25.3 mL of 0.50 M HCl?
55. How many milliliters of 1.25 M HCl must be added to 50.0 mL of
2.50 M LiOH to make a neutral solution?
56. What is a buffer? What do they consist of?
A buffer is a solution that resists a change in pH. Buffers
contain a weak acid and its conjugate base.
57. For the following questions label each part of the following chemical
reactions with acid (A), base (B), conjugate acid (CA) and conjugate
base (CB).
UNIT THIRTEEN:
EQUILIBRIUM
58. Terms
• Activation Energy - amount of energy that must be absorbed by
reactants before a reaction can occur
• Catalyst - a substance that increases the rate of a chemical reaction
but is not consumed by the reaction
• Equilibrium - when forward and reverse reactions are occurring under
the same conditions and at the same rate
• Enthalpy - the energy difference between the reactants and products;
another word for energy
• Entropy - randomness, disorder
• Equilibrium constant - product of concentrations of products divided by
the product of the concentrations of reactants
• Free Energy - the net balance or difference between energy (enthalpy)
and entropy
• Le Châtelier’s Principle - when a chemical system at equilibrium is
disturbed, the system shifts in a direction that minimizes the
disturbance
• Reversible Reaction - a reaction where the products can re-form
reactants
• Solubility Product Constant - is an equilibrium expression for a
chemical equation that represents the dissolving of an ionic
compound
• Spontaneous - happens on its own, does not have to be forces
• Static - at rest; not changing
• Steady State - constant change but unlike equilibrium there is no
reverse reaction
• Thermodynamic - the study of energy relationships
59. Describe each of the following situations as either “equilibrium” or “steady
state”.
a. Lake Wissota Dam and the water behind it. The level in the lake is
constant.
Steady State
b. The liquid in a sealed container and the vapor above it at constant
temperature
Equilibrium
c. The solid solute remaining in a sealed container of solution
Equilibrium
60. For the reaction below, Keq = 0.39. If the equilibrium concentration of
HC2H3O2 is 2.3 M, what are the equilibrium concentrations for the two ions
below?
HC2H3O2 (aq) + H2O (l) ⇌ H3O+(aq) + C2H3O2 – (aq)
61. Calculate the value of Keq for the reaction below, using the given equilibrium
concentrations:
CH4 (g) + 3 Br2 (l) ⇌ CHBr3 (aq) + 3 HBr (aq)
62. What causes a system at equilibrium to shift?
Stress
63. Methyl alcohol (“methanol”) is made according to the
following net equation. Predict the effect on the equilibrium
system (“shift left”, “shift right”, or “no effect”).
CO (g) + 2 H2 (g) ⇌ CH3OH (g) + heat
a. CO is added to the container:
b. The volume is decreased:
c. A catalyst is added:
d. Temperature is increased:
e. Pressure is increased:
f. Hydrogen gas is removed:
RIGHT
RIGHT
NO EFFECT
LEFT
RIGHT
LEFT
64. Write the complete equation for Fe(OH)2 and write the Ksp
expression.
Fe(OH)2 → Fe2+ + 2OHKsp = [Fe2+][OH-]2
•
UNIT FOURTEEN: REDOX
68. Assign an oxidation number to the specified atom in
each of the following examples.
71. Will the following redox reactions be spontaneous?
a. Pb(s) + Sn2+(aq) → Pb2+(aq) + Sn(s) NO
b. Pb(s) + Ag+(aq) → Pb2+(aq) + Ag(s)
YES
72. How can you pre-determine if a metal will dissolve in an
acid?
If the metal is below H2 on the activity series it will
not dissolve in the acid. If the metal is above H2 on the
activity series it will dissolve in an acid.
73. Calculate the cell voltage and give the cell notation for
electrochemical cells made from the following electrodes.
Ecell = 1.05v
Ni | Ni2+ || 2Ag+ | 2Ag
UNIT FIFTEEN: ORGANIC
80. Draw isomers of hexane.
UNIT SIXTEEN:
BIOCHEMISTRY
93. Dogs may eat homework but do they digest it? Explain why.
Paper is made of cellulose which is indigestible so the dog cannot digest
the paper.
94. The insulin protein contains 51 amino acids. How many DNA base pairs are
required to code for all the amino acids in insulin?
153
95. When wet hair is put into curlers and allowed to dry, the hair tends to retain the
shape of the curler. Why?
When hair dries the hydrogen bonds stay in place from when wet,
therefore keeping the shape of the curler.
96. Why is shape so important in the functions of proteins?
The shape determines the function. Heat can denature and change
shape therefore change function.
97. Name three lipids of the cell and describe their function.
Answer will vary
98. What are the different structures of proteins? How would you identify these
structures?
Primary – amino acid sequence
Secondary – short range repeating pattern along protein chain
Tertiary – large scale bends and folds
Quaternary – arrangement of chains in protein
99. How are alpha-helix proteins different from beta-pleated sheet proteins?
Alpha helix is a coil held together through hydrogen bonds of NH and CO.
The side chains extend outward.
For beta-pleated sheets the peptide backbones of the chain interacts
through hydrogen bonds. The side chains extend above and below.
100. The anti-HIV medicine AZT introduces a fake thymine molecule into the cell.
When a virus attempts to use the fake nucleotide in replicating its DNA, the fake
nucleotide doesn’t work. Below are diagrams of the fake thymine on the left and the
real thymine on the right. How does AZT keep the virus from replicating when the
virus tries to incorporate the fake thymine into its DNA?
Azidothymidine and thymine have the same shape but slightly different
structures. Thymine has an OH group and azidothymidine has an amine group (N=N=NH). The different structure has a different function.
Final Schedule
Friday:
1st Hour
2nd Hour
Lunch
3rd Hour
8:35 – 10:30
10:40 – 12:35
12:35 – 1:30
1:35 – 3:35
Monday:
4th Hour
5th Hour
Lunch
6th Hour
8:35 – 10:30
10:40 – 12:35
12:35 – 1:30
1:35 – 3:35