Download Answers to 2017 Chemistry Exam Review Compounds and

Answers to 2017 Chemistry Exam Review
Compounds and Reactions
1. 7 diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2
2. ionic compounds: metal cation – nonmetal anion; covalent – nonmetals
3. Use octet rule to share enough electrons for each atom to have 8 valence electrons.
4. polar – unequal sharing: H :Cl
nonpolar – equal sharing: H : H or Cl : Cl
3+
25. example: Al and O : Al2O3
6. 1: mono-, 2: di-, 3: tri-, 4: tetra-, 5: penta-, 6: hexa-, 7: hepta-, 8: octa-, 9: nona-, 10: deca7. ionic – use lists on back of PT and balance charge with subscripts; covalent – use prefixes
8. left-side substances = reactants, right-side substances = products
9. 5 types of reaction: synthesis (1 product), decomposition (1 reactant), single replacement
(element + compound on both sides), double replacement (2 ionic compounds on both
sides), combustion (fuel + O2 = CO2 + H2O)
10. Count atoms of each element and use coefficients, as needed, to equalize #atoms on each
side.
11. s = solid, l = liquid, g = gas, aq = aqueous (dissolved in water)
Stoichiometry
12. mole = Avogadro’s number = 6.02 x 1023; molar mass = mass of 1 mole = atomic mass in
PT in grams; molar volume of gas at STP = 22.4 L
13. Use subscripts to determine #atoms of each element: in Ca3(PO4)2: 3 Ca, 2 P, 8 O
14. Multiply molar mass of each element by subscript #s then add up total: In Ca3(PO4)2 = (3
x 40.08) + (2 x 30.97) + (8 x 16.00) = 310.18 g
15. %composition = element mass / compound mass x 100: In Ca3(PO4)2 = (3 x 40.08) /
310.18 = 38.8% Ca; (2 x 30.97) / 310.18 = 20.0 %; (8 x 16.00) / 310.18 = 41.2%
16. Stoichiometry is the study of amounts in chemical reactions.
17. #mol x g/molar mass = #g
#g x 1 mol/molar mass = #mol
#mol x Av. # / 1 mol = #atoms/molecules #atoms/molecules x 1 mol/Av # = #mol
#mole x 22.4 L / 1 mol = #L gas at STP
#L gas at STP x 1 mol / 22.4 L = #mol
18. In particle diagrams for reactions, use a different symbol for each element and show the
particles attached for compounds and separate for elements. Make sure the same total
number of particles exists for each element on each side of the equation.
19. Use coefficients from the balanced equation to convert from moles of one substance to
moles of another substance: #mol A x coeff. B / coeff. A = #mol B
20. To convert from grams of one substance to grams of another substance in a balanced
equation, multiply the coefficient times the molar mass of each substance then multiply
the known mass by the ratio of unknown/known to get unknown mass:
#g A x coeff B x molar mass / coeff A x molar mass = #g B
21. %yield = (actual yield / theoretical yield) x 100
Thermochemistry
22. System is whatever you are studying; surroundings is whatever is outside what you’re
studying. Energy is constant for system plus surroundings.
23. Heat is the movement of thermal energy.
24. Heat moves spontaneously from hot to cold. This is because the hotter (faster) molecules
speed up the colder (slower) molecules. Heat stops moving when the hot and cold
molecules reach the same temperature.
25. A refrigerator does work to move heat from cold to hot.
26. Endothermic is heat moving into the system while exothermic is heat moving out of the
system.
27. Specific heat is the resistance to change in temperature. Water has a high specific heat
(does not change temperature easily). Metals have low specific heat and change
temperature quickly.
28. Metals feel cool b/c they conduct heat away from your skin.
29. It is good that your body and Earth contain so much water so that temperature does not
change easily.
30. Enthalpy is chemical energy. In an exothermic process the change of enthalpy ( H) is
negative b/c the products have less enthalpy than the reactants. In an endothermic process
the change of enthalpy ( H) is positive b/c the products have more enthalpy than the
reactants.
31. Breaking bonds increases enthalpy; making bonds decreases enthalpy.
32. Enthalpy graph for endo- and exothermic reactions:
33. A calorimeter measures heat of combustion by measuring how much heat goes from the
burning substance into water using the equation Q = mC T for water.
34. We burned the fuel to heat the water then calculated the heat of combustion by
multiplying mass times specific heat times temperature rise for the water then dividing by
the mass of fuel that burned.
35. energy/mol of fuel = energy/g x molar mass of fuel; energy/g = energy/mol / molar mass
of fuel; energy per mol > energy/g b/c mol > g
36. The chemical formula with more bonds will have a greater energy per mole b/c the
energy is in the bonds.
Solutions
37. A solution is a solute dissolved in a solvent. The solvent for aqueous solutions is water.
38. Water is a bent molecule with the two hydrogen atoms (electronegativity 2.1) forming the
more positive pole and an oxygen atom (electronegativity 3.5) forming the more negative
pole. The shared electrons are closer to the more electronegative oxygen atom.
39. The water molecules “cling” to each other b/c the more electropositive H atoms of one
water molecule are attracted the more electronegative oxygen atoms of the next water
molecule.
40. The more electropositive H atoms are attracted to the Cl- anions and the more
electronegative O atoms are attracted to the Na+ cations.
41. Water tends to dissolve other polar substances since its positive pole will be attracted to
the negative pole of the other substance and vice versa. It tends not to dissolve nonpolar
substances b/c there are no poles for water’s poles to be attracted to.
42. Ca(NO3)2 Ca2+ + 2 NO3143. A precipitate is an undissolved solid that results from a double replacement reaction
between two soluble compounds.
44. example: 2 AgNO3 (aq) + MgCl2 Mg(NO3)2 (aq) + 2 AgCl (s)
complete ion eq: 2 Ag+ + 2 NO3- + Mg2+ + 2 Cl- Mg2+ + 2 NO3- + 2 AgCl
(Mg2+ & NO3- = spectator ions)
net ionic eq: 2 Ag+ + 2 Cl- 2 AgCl
45. Concentration of a solution can be increased either by (1) adding more solute or (2)
letting the solvent evaporate and can be decreased by adding more solvent.
46. Molarity is moles solute per liter solution. It is represented by brackets [substance].
47. To convert from ml to L shift your decimal place 3 places to the left.
48. To find molarity, divide #mol/#L. To find #mol, multiply molarity x #L. To find #L,
divide #mol by molarity.
Acids and Bases
49. An acid is a proton (H+) donor. A base is a proton (H+) acceptor.
50. An H+ ion is a proton (a hydrogen atom without its electron). H+ + H2O = H3O+.
51. A hydrogen is ionizable if it is bonded to an element with high electronegativity. This
pulls the shared electron away from the hydrogen atom.
52. dissociation of HCl in water: HCl + H2O = H3O+ + Cl53. H3O+ is a hydronium ion. OH- is a hydroxide ion. Acids are associated with hydronium
and bases with hydroxide ions.
54. Equilibrium is a reversible reaction where the rate of the forward and reverse reactions is
the same.
55. Le Chatelier’s principle says that, if an equilibrium reaction is stressed by the addition or
reduction of a substance, the equilibrium will shift to relieve the stress. For example, if
some product is removed, the equilibrium will shift to the right to replace the product. Or
if some product is added, the equilibrium will shift to the left to remove some of the
added product.
56. The conjugate acid is the acid on the product side that results from the base on the
reactant side gaining a proton (H+). The conjugate base is the base on the product side
that results from the acid on the reactant side losing its proton (H+). The conjugates look
like their counterparts on the reactant side – just with or without the proton (H+).
example: HCl + H2O = H3O+ + Clacid
base
conj acid conj base
57. Acids or bases differ from their conjugates by a proton (H+).
58. Amphoteric is a chemical that can act as an acid (proton donor) or a base (proton
acceptor).
For example, HCO3- can act as an acid by donating H+: HCO3- + H2O = H3O+ + CO32- .
Or it can act as a base by accepting a proton: HCO3- + H2O = H2CO3 + OH- .
(Notice that water is also amphoteric – acting as a base in the first example and an acid in
the second.)
59. Kw = [H3O+][OH-] = 1.0 x 10-14 This is a very small number, meaning water rarely selfionizes.
60. [H+] = 1.0 x 10-14 / [OH-] or [OH-] = 1.0 x 10-14 / [H+]
61. pH = the negative power of ten for [H+] or –log [H+]
[H+] = 2nd log –pH
62. For pure water pH = 7 and [H+] = 1.0 x 10-7 M
63. pH + pOH =14; so pOH = 14 – pH or pH = 14 – pOH
64. A pH unit is a power of ten for [H+]. An increase of one pH unit is a 1/10th decrease in
acidity; a decrease of one pH is a 10x increase in acidity.
65. Almost 100% of a strong acid ionizes, as indicated by a forward arrow, while only a
small fraction of a weak acid ionizes, indicated by a double arrow.
66. A strong acid ionizes almost entirely. A concentrated acid has a lot of acid per volume of
solution.
67. concentrated strong
dilute strong
concentrated weak
dilute weak
68. Ka measures strength of an acid. Ka = [H3O+][A-] / [HA]
69. ex: If [H3O+] = [A-] = 0.0001 M and [HA] = 0.1 M, Ka = (0.0001)2 / 0.1 = 1.0 x 10-7
70. ex: If Ka = 5.2 x 10-8 and [HA] = 0.1 M, 0.1 M = x2 / 0.1; x = 7.2 x 10-5 M
71. acid + base = salt + water; net ionic equation = H+ + OH- = H2O
72. Titration is finding the amount of an acid by neutralizing it with a known amount of base,
or vice versa.
73. (Macid)(Lacid) = (Mbase)(Lbase) – If you know 3 of the 4, you can calculate the 4th.
Gases
74. Gases are much less dense than liquids or solids b/c their molecules are spread out.
75. Temperature relates to the kinetic energy or speed of molecules. Pressure relates to how
often and how hard the molecules hit each other or other things.
76. Atmospheric pressure depends on the weight of air above = about 15 psi = 1 atm at sea
level. It decreases altitude increases b/c of less weight of air above.
77. Decrease volume (as in a capped syringe), increase gas (as in pumping air into a rigid
container), or increase temperature in a rigid container.
78. Draw larger circles for greater volume, more dots for more gas, and longer arrows for
higher temperature.
79. Temperature must be measured in Kelvins (K) for gas laws. This is because you can’t
have negative values for things like volume and b/c the values are proportional to the
lowest temperature being absolute zero (when there is no molecular motion).
80. When temperature rises, volume rises in a balloon or pressure rises in a rigid container.
This is b/c the gas molecules move faster and hit the sides of the container harder and
more often.
81. When external pressure on a balloon decreases, its volume increases b/c there are fewer
molecules hitting the outside of the balloon.
82. If you boil water in an open soda can, you force out the air so that it will implode when
you put it upside down in water. The outside air pressure crushes it.
83. A sub rises when enough air is added to make it less dense than the surrounding water. It
sinks when enough water is added to make it more dense than the surrounding water.
84. To get dry pressure, subtract water-vapor pressure from the total air pressure.
85. Remember to convert to Kelvin temperatures. It also helps to label each measurement
with a letter (for example, “P1” for the first pressure). Also R = 0.082.
Electrochemistry
86. Remember that the oxidation number is -2 for oxygen, +1 for hydrogen, and the charge
on back of PT for metals. Remember also that neutral elements and compounds have
oxidation number 0 and that polyatomic ions (such as CO32-) have a total oxidation
number equal to their charge. For example, in CO32-: C + 3(-2) = -2; C – 6 = -2; C = +4.
87. In oxidation an element loses electrons for a larger oxidation number. In reduction an
element gains electrons for a smaller oxidation number. In a redox reaction (and
oxidation always goes with reduction), electrons move from the oxidized to the reduced
element.
88. In a redox reaction, the element whose oxidation number increases is oxidized; the
element whose oxidation number decreases is reduced.
89. The more active metal is the one that is oxidized – it more readily loses its electrons to
become a cation.
90. Oxidation occurs at the anode. (Remember “an ox.”). Reduction occurs at the cathode.
(Remember “red cat.”)
91. oxidation at anode: M M+ + ereduction at cathode: M+ + e- = M
The salt bridge connects the anode and cathode reactions and keeps each solution neutral.
92. A voltaic cell converts chemical into electrical energy. An electrolytic cell uses electrical
energy to cause a chemical change.
93. In electroplating a metal cation is reduced on a surface at the cathode.