3Chemistry Final’s Review
The review is a tool, not an answer sheet to be memorized. You will find everything we have discussed in class below. Use your notes and credible
internet sources (outside of class) to help you answer the questions, concepts, or comments in need of elaboration. You should look at this 1-3 times
a day, go to bed with it, eat your meals with it, care for it, for it will guide you to the grade you want on the final. The more you practice and review
the better you will do.
Standard 1.1 Analyze the structure of atoms and ions.
1. Describe the location, relative charge, purpose, and relative mass of all three subatomic particles.
a. Proton – Positive (+1), found in nucleus, mass of 1, defines element and attracts electrons
b. Neutron – Neutral (+/- 0), found in nucleus, mass of 1, stabilize nucleus and define isotope
c. Electron – Negative (-1), found in electron cloud surrounding nucleus, mass negligible, responsible for
reactions
2. A = mass number, Z = atomic number – Ex. AZX
3. What are the two ways isotopes can be expressed (written)? Give an example of each.
a. Carbon-14
b. C-14
4. What is the difference between C-12 and C-14?
a. C-14 has two more neutrons than C-12 (They are isotopes)
5. Determine the number of each subatomic particle for the following isotopes:
Symbol
Atomic number
Mass number
Number of
Number of
Number of
Protons
Electrons
Neutrons
6.
7.
8.
9.
10.
11.
12.
13.
14.
23Na
11
23
11
11
12
K
19
40
19
19
21
O2-
8
16
8
10
8
N3-
11
23
11
14
12
Cl1-
17
35
17
18
18
S
16
32
16
16
16
What does amu stand for?
a. Atomic Mass Unit
Explain why atoms have different isotopes. In other words, how is it that helium can exist in three different forms?
a. Each has a different amount of neutrons in their nuclei
How do you find average atomic mass? (What is the formula?)
a. AAM = (mass of isotope x %relative abundance) + (mass of isotope x %relative abundance)…
The natural abundance for boron isotopes is: 19.9% 10B (10.013 amu) and 80.1% 11B (11.009 amu). Calculate the atomic weight of boron.
a. (10.013 amu x .199) + (11.009 amu x .801) = 10.810796 amu
Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the
average atomic mass of rubidium?
a. (85 amu x .722) + (87 amu x .278) = 85.556 amu
Uranium has three common isotopes. If the abundance of 234U is 0.01%, the abundance of 235U is 0.71%, and the abundance of 238U is
99.28%, what is the average atomic mass of uranium?
a. (234 x .0001) + (235 x .0071) + (238 x .9928) = 237.9783 amu
Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of
titanium?
a. (46 x .08) + (47 x .078) + (48 x .734) + (49 x .055) + (50 x .053) = 47.923
Copper exists as two isotopes: You have 30 Copper-63 isotopes and 470 Copper-65 isotopes. 63Cu (62.9298 amu) and 65Cu (64.9278 amu).
What are the percent abundances of the isotopes?
a. AAM(find in periodic table) = (given isotope1 amu) x (𝑥) + (given isotope1 amu) x (1 − 𝑥)
b. Remember: part/whole
c. Cu-63 = 6%
d. Cu-65 = 94%
What’s the difference between how we draw an atom and its theoretical shape?
a. In our drawing of an atom which model do we use? (scientist’s name)
b. Bohr Model has PELs represented as rings like a solar system’s orbits of its planets
c. Theoretical shape is an electron cloud or probability cloud that shows the probability of encountering an electron in a certain area
15. Describe the concepts of excited and ground state of electrons in an atom.
a. Ground state is an energy state where the electron has no additional energy. It is “happy and content”.
b. Excited state is an energy state where an electron has additional energy that it must get rid of in the form of
radiation. It is “overjoyed”.
16. How can an electron move from a ground state to an excited state?
a. Describe if the movement is exothermic, or endothermic and why.
b. Electrons become excited by gaining energy (Endothermic – absorbing energy)
17. What happens when an electron moves from an excited state to a ground state?
a. Describe if the movement is exothermic, or endothermic and why.
b. Electrons move to ground state by losing energy (Exothermic – releasing energy)
18. Energy, in an atom, exists as a discrete unit called, what?
a. Quanta
19. What is electromagnetic radiation? Physically speaking, how is electromagnetic radiation given off?
a. “EM radiation is created when an atomic particle, such as an electron, is accelerated by an electric field,
causing it to move. The movement produces oscillating electric and magnetic fields, which travel at right
angles to each other in a bundle of light energy called a photon.” (LiveScience)
20. Describe the relationship between wavelength and frequency.
a. The higher the frequency, the shorter the wavelength
21. How do the following compare: wavelength and energy, frequency and energy, and amplitude and energy.
a. More energy = shorter wavelength
b. More energy = higher frequency
c. More energy = greater amplitude
22. Describe the relationship between wavelength and color.
a. Longer to Shorter wavelength
b. Red, orange, yellow, green, blue, indigo, violet
23. Electrons “circle” the nucleus in energy ranges called what?
a. PELs
24. How do electrons move from one PEL (answer from previous question) to another by doing what?
a. Gaining or losing energy
25. Describe the wave/particle duality of electrons.
a. Probability waves describe the possibility of finding an electron in a given location
26. Draw and describe the symbols for alpha particles, beta particles, and gamma radiation.
a. Alpha – α
b. Beta – β
c. Gamma – γ
27. Describe each particle/radiation’s strength.
a. Alpha – big and slow, weak but lethal upon contact or if ingested
b. Beta – stray electron or positron, stronger but can be shielded by aluminum foil
c. Gamma – Strongest, very lethal, can cause severe burns. Shielded by inches of lead.
28.
29. What is half-life?
a. Time it takes for half of a sample to decay
b. A = A0 x (½)t/h
c.
30. Compare radioactive decay with fission and fusion.
a. Radioactive decay is in the form of fission where nuclei separate into lighter nuclei. Releases large amounts of
energy.
b. Fusion is when lighter nuclei fuse to create heavier nuclei. Releases tremendous amounts of energy.
Standard 1.2 Understand the bonding that occurs in simple compounds in terms of bond type, strength, and properties.
31. Describe metallic bonds.
a. Metals share free electrons between one another. This gives the metals good conductivity and structure.
32. What is the difference between a cation and anion?
a. Cations are positive+
b. Anions are negative33. What type of bond is created using a cation and anion?
a. Ionic bond
34. What type of bond is a sharing of electrons?
a. Covalent bond
35. What type of bond is a melding of electrons where they act as if in a sea, interacting with one-another.
a. Metallic bond
36. What does the term oxidation state refer to? When do you use oxidation states? How do you correctly apply them to a formula?
a. Oxidation state is how many electrons an atom needs or has to give away to complete an octet
b. Used when creating chemical compounds
c. They are superscripts first. Then they are crisscrossed down to the subscript. Reduce if possible
37. Describe the method of writing, in symbol form, Lithium Oxide, Iron (II) Sulfide, Aluminum Phosphide, Magnesium Chloride, and Manganese
(IV) Oxide. They all differ so please write the steps for each
a. Li1 + O2 → Li2O
b. Fe2 + S2 → FeS
c. Al3 + P3 → AlP
d. Mg2 + Cl1 → MgCl2
e. Mn4 + O2 → Mn2O
38. Name the following: MgO, NaCl, NO2, Fe2O3, NaBr, CaO, Li2S, MgBr2, Be(OH)2, SO3, N2S, PH3, BF3, P2Br4, CO, SiO2, SF6, NH3, NO2
a. Magnesium Oxide
k. Dinitrogen Sulfide
b. Sodium Chloride
l.
Phosphorus Trihydrogen
c. Nitrogen Dioxide
m. Boron Trifluoride
d. Iron (III) Trioxide
n. Diphosphorus Tetrabromide
e. Sodium Bromide
o. Carbon Monoxide
f. Calcium Oxide
p. Silicon Dioxide
g. Dilithium Sulfide
q. Sulfur Hexafluoride
h. Magnesium Dibromide
r. Nitrogen Trihydrogen
i.
Beryllium Hydroxide
s. Nitrogen Dioxide
j.
Sulfur Trioxide or Sulfite
39. Use symbols to name the following: nitrogen trichloride, boron carbide, dinitrogen trioxide, phosphorus pentafluoride, methane, sulfur
dibromide, diboron tetrahydride, oxygen difluoride, carbon disulfide, nitrogen monoxide, potassium iodide, magnesium oxide, aluminum
chloride, sodium nitrate, calcium carbonate, lithium sulfate, beryllium phosphide, magnesium hydroxide, sodium phosphate, aluminum
carbonate, calcium chloride, sodium cyanide, aluminum oxide, magnesium acetate, ammonium chloride.
a. NCl3
b. B4C3
c. N2O3
d. PF5
e. CH4
f. SBr2
g. B2H4
h. OF2
i.
Li3SO4
j.
Be3P2
k. Mg(OH)2
l.
Na3P
m. Al2(CO3)3
n. CaCl2
o. NaCN
p. Al2O3
q. Mg(C2H3O2)2
r. NH4Cl
40. Draw the Lewis Dot Diagram for the following: PBr3, N2H2, CH3OH, NO2-1, C2H4, H2, O2, Br2, F2, I2, N2, Cl2, CH4, CS2, N2O, CCl4
a. *Attached
41. In terms of electronegativity values, what would covalent and polar covalent bond exist in? (number value)
a. 0.5-1.7
42. In terms of electronegativity values, what would an ionic bond exist in? (number value)
a. 1.7+
43. Fill in the following table. Only use covalent, polar covalent, and ionic
44. Why are hydrogen bonds stronger than dipole-dipole forces?
a. Hydrogen atoms are smaller and can get closer to their counterpart
45. Why are dipole-dipole forces stronger than dispersion forces?
a. Dipole-dipole forces are more stable than a dispersion force
46. Describe the relationship between bond energy and length of single, double, and triple bonds.
a. Triple bonds are shorter than single bonds because they require more energy to break, thus are stronger
47. Write, with charges, these polyatomic ions: nitrate, sulfate, carbonate, acetate, and ammonium.
a. NO3b. SO42c. CO32d. C2H3O248. Write the names for these common acids: HCl, HNO3, H2SO4, HC2H3O2.
a. Hydrochloric Acid
b. Nitric Acid
c. Sulfuric Acid
d. Acetic Acid
49. What are the characteristics of ionic bonds in terms of: MP, BP, electrical conductivity, and malleability?
a. High MP and BP, low conductivity, not malleable
50. What are the characteristics of covalent bonds in terms of: MP, BP, electrical conductivity, and polarity?
a. Low MP and BP, low conductivity, polar
51. What are the characteristics of metallic bonds in terms of: MP, BP, conductivity, malleability, ductility, and luster?
a. High MP and BP, high conductivity, high malleability, ductility, and luster
52. Describe VSEPR theory.
a. Valence Shell Electron Pair Repulsion Theory
b. Predicts shapes of compounds. Electrons repel each other so atoms around a central atom will be spread out
as much as possible.
53. For each of the following compounds, determine the bond angles, and molecular shapes: carbon tetrachloride, BH 3, silicon disulfide, C2H2,
PF3, PF6-, PCl5, H2O, nitrogen trihydride
a. CCl4 – Tetrahedral – 109.5°
b. BH3 – Trigonal Planar – 120°
c. SiS2 – Linear – 180°
d. C2H2 – Linear – 180°
e. PF3 – Trigonal Pyramidal – ?°
f. PF6 – Octahedral – 90°
g. PCl5 – Trigonal Bipyramidal
h. H2O – Bent – 104.5°
54. How does bond angle relate polarity?
a. Depending on the shape of the compound, one can determine if it is a dipole or not.
55. What is a dipole?
a. Compound that has two poles. Can be divided into a positive and negative end.
56. Determine if following is either polar or non-polar by using the correct VSEPR drawing and using a dipole(s): H2, H2O, CH4, NH3, HF, C2H2,
CH2Cl2, N2, CH3OH, H2S, PCl3, CO2, and C6H6.
a. *Attached
57. Describe these macromolecules and their unique properties: water (ice), graphite/diamond, polymers (PVC, nylon), and proteins (hair,
DNA).
a. Ice – forms crystal arrangements that trap oxygen inside, making it less dense than its liquid form.
b. Graphite/diamond – Graphite, because it has a free electron, conducts electricity. Diamonds form crystal
arrangements that make them some of the hardest substances known to man.
c. Polymers – long chains of carbons and halogens that make these substances inert, strong, and resistant to
heat and corrosion.
d. Proteins – serve a specific purpose. Can be “programmed” and can retain information. All are essential to life.
Standard 1.3 Understand the physical and chemical properties of atoms based on their position on the Periodic Table.
58. Name and describe the properties, valence electrons, oxidation state and reactivity of the groups on the periodic table.
a. Alkali Metals – Highly reactive, 1 valence electron, oxidation state of -1,
b. Alkaline Earth Metals – Not as reactive as Alkali Metals, 2 valence electrons, oxidation state of -2
c. Transition Metals – More reactive to the left, unknown oxidation state or valence electrons.
d. Group 13 – Moderately reactive, 3 valence electrons, oxidation state of -3
e. Group 14 – Moderately reactive, 4 valence electrons, oxidation state of +/- 4
f. Group 15 – More reactive, 5 valence electrons, oxidation state of +3
g. Group 16 – More reactive, 6 valence electrons, oxidation state of +2
h. Group 17 – Most reactive, 7 valence electrons, oxidation state of +1
i.
Group 18 – Inert, 8 valence electrons, oxidation state irrelevant/0
59. Reactivity increases as you go top right on the periodic table and decreases as you go bottom center on the periodic table.
60. Identify the regions of the periodic table that are metals, nonmetals, and metalloids.
a. Metals – Entire left side
b. Nonmetals – top diagonal right side
c. Metalloids – Staircase starting above but excluding aluminum
61. Identify which blocks are s, d, p, and f.
a. s – Alkali Metals, Alkaline Earth Metals, Hydrogen, and Helium
b. p – Group 13 to 18
c. d – Transition Metals
d. f – Rare Earth Metals
62. Rank the following elements by increasing atomic radius: Carbon, Aluminum, Oxygen, and Potassium.
a. Oxygen, Carbon, Aluminum, Potassium
63. Describe the general trends within groups and periods for electronegativity.
a. Electronegativity increases going from left to right and bottom to top.
64. Rank the following elements by increasing electronegativity: Sulfur, Oxygen, Neon, and Aluminum.
a. Neon, Aluminum, Sulfur, Oxygen
65. Describe the general trends within groups and periods for ionization energy.
a. Ionization energy is like electronegativity
b. Ionization energy is how much energy is needed to rip off an electron
66. Why does fluorine have higher ionization energy than iodine?
a. Fluorine’s electrons are closer to its nucleus and carry a higher effective nuclear charge
67. Indicate whether the following properties increase or decrease from left to right on the periodic table and WHY.
a. Atomic radius – Decrease – Electrons are pulled closer to the nucleus
b. Ionization energy – Increase – Harder to rip off an electron
c. Electronegativity – Increase – Atoms want an electron more badly
68. Write the electron configuration for the following elements: Li, O, P, Xe, Cu, and K.
a. *Attached
69. Use the noble gas configuration to write the electron configuration for the following: Sn, Cr, Mn, Ar, and Cs.
a. *Attached
70. Use orbital diagram to indicate the following elements: Si, Na, Cl, Zn, Sr, F, and C.
a. *Attached
71. Identify the oxidation numbers for all groups excluding the transition metals.
a. Group 1 – -1
b. Group 2 – -2
c. Group 13 – -3
d. Group 14 – +/-4
e. Group 15 – +3
f. Group 16 – +2
g. Group 17 – +1
h. Group 18 – 0
72. How does size lead to reactivity, gaining, or losing electrons.
a. The bigger the atom, its outer electrons are easier to rip off or gain. More reactive.
Standard 2.1 Understand the relationship among pressure, temperature, volume, and phase.
73. Explain how the energy (kinetic and potential) of particles of a substance changes when heater, cooled, or changing phase.
a. Kinetic energy rises the further up on the heating/cooling curve. When higher up, there is more potential
energy. When changing phase, kinetic energy rises but temperature stays the same.
74. Identify pressure as well as temperature as a determining factor for phase of matter.
a. Higher pressures raise melting/boiling points
75. Compare and contrast Joule, Celsius, and Kelvin.
a. Joule = Energy
b. Celsius = Measures kinetic energy based on freezing point of water
c. Kelvin = same scale as Celsius, but is absolute; based on zero energy
76. How does water stay as a vapor?
a. Energy is still being added or removed, but temperature stays the same.
77.
78. The above diagram is a heating/cooling curve. If you have any questions about the graph please ask me.
a. ok
79. Describe whether heat is being absorbed or released for water to become ice, ice to become vapor, vapor to become water, and ice to
become water.
a. Released
b. Absorbed
c. Released
d. Absorbed
80.
81. Describe the following graphs.
a. Phase diagrams of two different compounds.
82. What is a triple point for a material, in terms of a phase diagram?
a. A point where all three phases are indistinguishable from one another.
83. For a closed system, energy is neither lost nor gained only transferred between components of the system.
a. True
84. Identify the following equations and state their purpose for water: q = mCp∆T, q = mHf, and q = mHv.
a. q = mCp∆T is energy change with only one phase
i. q = energy in Joules
ii. m = mass in grams
iii. C = specific heat capacity
iv. ∆T = change in temperature (T1 – T0)
b. Energy required for Phase change
i. q = mHf is heat of fusion
ii. q = mHv is heat of Vaporization
85. Identify the characteristics of ideal gases.
a. Many identical molecules
b. Molecules occupy waaaay less space than the gas itself
c. Molecules move randomly
d. Molecules only experience force during collisions, which are elastic
86. 1 mole of anything is how many particles or molecules?
a. 6.02 x 1023
87. 1 mole of Oxygen, Potassium, or Iodine weighs how much?
a. 1 mol O2 = 15.9994g O2
b. 1 mol K = 39.0983g K
c. 1 mol I2 = 126. 90447g I2
88. 1 mole of Neon gas occupies how much space at STP?
a. 22.4 L = 1 mol
89. What does STP stand for?
a. Standard temperature and pressure
90. 1 atm = 101.325 kPa = 760 torr = 760 mm Hg
a. True
91. How many atm of pressure are there in 122.3 kPa?
a. 1.2 atm
92. What is the pressure of a gas in atm, if its pressure is 305 kPa?
a. 3.01 atm
93. Use equations 3-6 on the left hand side of your formulas cheat-sheet to answer the following:
a. A sample of oxygen gas occupies a volume of 250 mL at a pressure of 740 torr. What volume will it occupy at 800 torr?
i. 231.25 mL
b. A sample of nitrogen occupies a volume of 250 mL at 25 degrees Celsius. What will it occupy at 95 degrees Celsius (use v/t = v/t)?
i. 950 mL
c. A gas in a rigid container is found to exert 97.0 kPa at 25.0 degrees Celsius. What would the temperature be at STP (use p/t = p/t)?
i. 0 Celsius (If talking about STP on Temperature)
ii. 26.11 Celsius (If STP is only Pressure)
d. The partial pressure of F2 is 300 torr in a mixture where the total pressure is 1.00 atm. What is the mole fraction of F2?
i. Pgas = (molgas/moltotal) x (Ptotal)
ii. 300 = (x) x (760)
iii. x = 15/38
94. Use the following R values to solve the subsequent questions:
a.
b.
c.
R = 0.08206 L atm/mol K
R = 62.4 L mm Hg/mol K
R = 8.314 L kPa/mol K
when P is in atm
when P is in mm Hg or torr
when P is in kPa
i. PV = nRT
ii. n = moles
iii. R = Universal Gas Constant (3 R values)
iv. Used to convert liters of gas to moles & vice-versa when not at STP
d. How many moles of oxygen will occupy a volume of 2.5 L at 1.2 atm and 25 degrees Celsius?
i. 1.46 mol O2
e. What volume will 2.0 moles of N2 occupy at 720 torr and 20 degrees Celsius?
i. 3.47 L
f. What pressure will be exerted by 25 g of CO2 at a temperature of 25 degrees Celsius and a volume of 500 mL?
i. 1772.34 mmHg
g. At what temperature will 5.00 g of Cl2 exert a pressure of 900 torr at a volume of 750 mL?
i. 153.4 Celsius
h. How many moles of N2 will occupy a volume of 347 mL at 6680 torr and 27 degrees Celsius?
i. 1.38 mol N2
95. Name two ways you can make water boil? Explain how this is possible?
a. Increase temperature - this will add kinetic energy to the molecules and allow them to break free And change
state.
b. Decrease pressure - this will create a void which the water molecules will want to fill; they will break free and
change state.
Standard 2.2 Analyze chemical reactions in terms of quantities, product formation, and energy.
96. Explain collision theory
a. Particles move randomly and collide with each other, gaining or losing no net energy in the process.
97. What is activation energy?
a. Energy required to start a reaction.
98. Name three ideas that can lower activation energy.
a. Using a catalyst
b. Increasing Temperature and pressure
c. Increasing surface area
99. What is a precipitate?
a. Solid that forms or precipitates from an aqueous solution
100. What rules do you follow to identify if a precipitate will form?
a. Solubility rules
101. Color change is a physical or chemical change?
a. Chemical in most cases
102. Color intensity is a result of dilution (typically the more diluted a solution the lighter the color)
a. True
103. ΔH = measure of a system’s enthalpy (sum of internal energy of a system plus the product of the system's volume and pressure)
104. What is the difference between an endothermic reaction and an exothermic reaction?
a. Endo = absorbs heat
b. Exo = releases heat
105. ΔH (+) = Endothermic - system absorbs energy
106. ΔH (-) = Exothermic - system releases energy
107. Write the complete, balanced equation for the following:
a. Zinc(II) + Hydrogen Chloride
i. Zn+2 + 2HCl → ZnCl2 + H2
b. Potassium Iodide + Lead(II) Nitrate
i. 2KI + Pb(NO3)2 → 2KNO3 + PbI2
c. Decomposition of water
i. 2H2O → 2H2 + O2
d. Burning of C3H8
i. C3H8 + 5O2→ 3CO2 + 4H2O
e. Synthesize bromine and magnesium
i. Br-1 + Mg+2 → MgBr2
108. acid-base neutralization is a double replacement reaction
a. True
109. What is a hydrocarbon?
a. A compound containing only Hydrogen and Carbon
110. What is the conservation of mass?
a. The total mass of the reactants = total mass of the products always
111. Activity series is used to determine if, what type of reaction will take place?
a. Single displacement
112. How will a single displacement reaction NOT occur?
a. If Active metal is lower than the compound’s metal in the Activity Series
113. How will a double displacement reaction NOT occur?
a. If both products are aqueous, there was no reaction because products are soluble in water
114. What reaction type are solubility rules used?
a. Double displacement
115. Take the reaction: NH3 + O2
NO + H2O.
a. In an experiment, 3.25 g of NH3 are allowed to react with 3.50 g of O2. Write a balanced equation.
i.
4NH3 + 5O2 → 4NO + 6H2O
b. Solve for the product NO. Which reactant is the limiting reagent?
i.
O2 - LR
c. How many grams of NO can possibly be formed?
i.
2.6g NO
d. How much of the excess reactant remains after the reaction?
i.
1.77g NH3
116. What is the empirical formula of a compound that contains 0.783 g of carbon, 0.196 g of hydrogen and 0.521 g of oxygen?
a. C2H6O
117. Determine the molecular formula of a compound with an empirical formula of NH2 and a formula mass of 32.06 g/mol.
a. N2H4
118. Find the percent composition of (NH4)2S.
a. N – 41.12%
b. H – 11.83%
c. S – 47.06%
119. Reaction rate is proportional to the number of effective collisions.
a. True
120. How does temperature (kinetic energy), concentration, and/or pressure affect the number of collisions for solid, liquid, and gases?
a. Higher temperature, concentration, and pressure will increase the number of collisions.
121. What is a catalyst?
a. Different compound or agent that lowers the activation energy but is not a part of the reaction product
122. Define chemical equilibrium: rates of forward and reverse reactions are the same. If more product is added, reactants will be formed until
there is an equal proportion, and vice-versa.
a. Write the following equation in Keq form: H2(g) + I2(g) ←> 2HI(g)
i.
b.
[𝐻2 ]+[𝐼2 ]
[𝐻𝐼]2
Write the following equation in Keq form: PCl5(g) ←> PCl3(g) + Cl2(g)
[𝑃𝐶𝑙3 ]+[𝐶𝑙2 ]
i.
[𝑃𝐶𝑙 ]
5
123. Define Le Chatelier's principle: when a system at equilibrium is subject to stress, the system will shift in the direction that relieves stress.
a. What happens if you add or remove a reactant or product?
i. System will shift to produce more of either, eventually reaching equilibrium again.
b. What happens if you add or remove heat?
i. System will shift toward the less-energetic side
c. What happens if you increase or decrease pressure?
i. System will shift toward the side with less pressure
124. How can you identify and acid or base solely on their formula?
a. Acids usually have H in front
b. But Acids will donate an H+ ion and Bases will accept it
125. What are the chemical properties of an acid, base?
a. Acids - corrosive, high H+ ion concentration
b. Bases - corrosive, counter acids, high OH- concentration
126. What is the difference between concentration and strength of acids and bases?
a. The higher H+ concentration for acids, the stronger they are
b. The Higher the OH- concentration for bases, the stronger they are
127. 4 compared to 5 on the pH scale is 10x as strong.
a. True because it is a negative logarithmic scale
128. pH = 4.16 [H+] = 6.92 x 10-5 pOH = 9.84 [OH-] = 1.45 x 10-10
129. pH = .97 [H+] = 1.06 x 10-1 M pOH = 13.02 [OH-] = 9.43 x 10-14
130. pH = 10.18 [H+] = 6.61 x 10-11 pOH = 3.82 [OH-] = 1.51 x 10-4
131. pH = 7.93 [H+] = 1.17 x 10-8 pOH = 6.07 [OH-] = 8.53 x 10-7 M
132. For any dilution problem use the equation M1V1 = M2V2
a. What mass of CH3OH is required to prepare 1.50 L of 3.00 M solution?
i. “M = Moles / liters
ii. 3.00 = moles /1.50
iii. 3.00 * 1.50 = moles
iv. 4.50 = moles
v. 4.50 moles of CH3OH * 32.03 grams / 1 mole = 144.14 grams of CH3OH”
vi. - Kristy
133. For any titration problem use the equation MaVa = MbVb
a. ok
134. What is the difference between electrolytic and non electrolytic solutions?
a. “Electrolytes are salts or molecules that ionize completely in solution. As a result, electrolyte solutions readily
conduct electricity. Nonelectrolytes do not dissociate into ions in solution; nonelectrolyte solutions do not,
therefore, conduct electricity.” - Boundless
135. Summarize all colligative properties (vapor pressure, boiling point, freezing point, and osmotic pressure)
a.
136. Use the graph to the left to answer the following questions.
a. ok
137. At what state is potassium iodide when there is 140 g of potassium iodide in
100 g of water at 10 degrees Celsius?
a. Saturated
138. Which compounds decrease in solubility with an increase in temperature?
a. HCl, NH3, SO2
139. How many grams of sodium chloride and ammonium chloride would I need in
100 g of water (in separate beakers) at 20 degrees Celsius, to reach the saturation
point?
a. NaCl - 37g
b. NH4Cl - 37g
140. Solubility is the attractiveness of a substance to another substance. The less
attracted the substance is the more insoluble it is.
a. True
141. Solve the following net ionic equations:
a. K3PO4 + Al(NO3)3 →
b. BeI2 + Cu2SO4 →
c. Ni(NO3)3 + KBr →
d. cobalt(III) bromide + potassium sulfide →
e. barium nitrate + ammonium phosphate →
f. calcium hydroxide + iron(III) chloride →
142. Consider the following unbalanced reaction: 2Al + 6HBr → 2AlBr3 + 3H2
a. When 3.22 moles of Al reacts with 4.96 moles of HBr, how many moles of H2 are formed?
i. 2.48 mol H2
b. What is the limiting reactant?
i. HBr
c. For the reactant in excess, how many moles are left over at the end of the reaction?
i. 0.83 mol Al
143. Consider the following unbalanced reaction: 3Si + 2N2 → Si3N4
a. When 21.44 moles of Si reacts with 17.62 moles of N2, how many moles of Si3N4 are formed?
i. 7.15 mol Si3N4
b. What is the limiting reactant?
i. Si
c. For the reactant in excess, how many moles are left over at the end of the reaction?
i. 7.15 mol N2
144. Consider the following unbalanced reaction: CuCl2 + 2KI → CuI + 2KCl + I2
a. When 0.56 moles of CuCl2 reacts with 0.64 moles of KI, how many moles of I2 are formed?
i. 0.32 mol I2
b. What is the limiting reactant?
i. KI
c. For the reactant in excess, how many moles are left over at the end of the reaction?
i. 0.24 mol CuCl2
145. Consider the following unbalanced reaction: 4FeS2 + 11O2 → 2Fe2O3 + 8SO2
a. When 26.62 moles of FeS2 reacts with 5.44 moles of O2, how many moles of SO2 are formed?
i. 3.96 mol SO2
b. What is the limiting reactant?
i. O2
c. For the reactant in excess, how many moles are left over at the end of the reaction?
i.
146. Consider the following unbalanced reaction: C3H8 + O2 → CO2 + H2O
a. If you start with 14.8 g of C3H8 and 3.44 g of O2, determine the limiting reagent.
i.
b. determine the number of moles of carbon dioxide produced.
i.
c. determine the number of grams of H2O produced
i.
d. determine the number of grams of excess reagent left
i.
147. Consider the following unbalanced reaction: Al2(SO3)3 + NaOH → Na2SO3 + Al(OH)3
a. If 10.0 g of Al2(SO3)3 is reacted with 10.0 g of NaOH, determine the limiting reagent.
i.
b. Determine the number of moles of Al(OH)3 produced.
i.