Chem IA
Name ___________________________ Blk _____
Final Exam Review Packet
Measurements/Unit Cancellation/Significant Figures
1. When measuring a cube with a ruler, you notice that an edge of the cube is between 2 and 3 centimeters. If
the ruler has marks every tenth of a centimeter, how many places after the decimal point do you need to have?
Answer: 2 places because you definitely know what the tenths place will be, and you have to estimate
a guess for the hundredths place.
2. How many significant figures will the answer to this problem – 200.45×.005500×100.0×10 – have?
Answer: 1 significant figure because the number 10 has only one significant figure. For
multiplication you use the number with the lowest amount of significant figures for the final amount of
significant figures.
3. How many inches are in 2.0 kilometers?
Answer: 2.0 km×(1000m)(100cm)( 1in ) = 79,000 in
(1km) (1m) (2.54cm)
4. How many millimeters are in a 1.000 mile?
Answer: 1.000 mi×(5280ft)(12in)(2.54cm)(10mm) = 1,609,000mm
(1mi) (1ft) (1in)
(1cm)
5. Lets say that your bedroom is 9 ft long by 13 ft wide by 9.5 ft high. If an air conditioner can exchange air
at a rate of 3 cubic yards per minute, how long will it take for the air conditioner to exchange all the air in
your room once (always taking in air that has not been exchanged yet). Make sure that your answer has
correct significant figures.
Anwer: 9 ft ×13 ft × 9.5 ft = 1000 cubic ft
3 cubic yards × (9 cubic ft/1 cubic yard) = 27 cubic feet
1000 cubic feet/27cubic feet = 40
It will take 40 minutes to exchange all the air in the room.
It is one significant figure because of the length of the room which is only one significant
figure
6.
How many significant figures should the answer to this equation be?
45g HCl
1 mole HCl
1 mole NaCl
58.44g NaCl
X
X
X
1
36.458g HCl
1 mole NaCl
1 mole HCl
Answer: 2 significant figures
7.
How many feet are in 2.5 miles?
Answer: 13200 feet
8.
How many significant figures should the answer to this equation be?
150.0 g H2O
+ 0.507 g salt
X g of solution
Answer: 4 sig. Figs because the answer cannot have more decimal places than the least
accurate measurement
9.
How many decimeters are there in 1 meter?
Answer: 10 decimeters in 1 meter
10.
What is the measurement of this? (DON’T FORGET TO USE 1/5 RULE)
Answer: 5.32 cm because the smallest unit is .1 of a cm so if that is divided by five you get .02 cm.
Naming and Reaction types
KEY TERMS
Combination/Synthesis: A reaction in which substances combine to form a new compound.
1. element + element => binary compound (A + B => AB)
2. metal oxide + water => base
3. Nonmetal oxide + water => acid
4. Metal oxide + nonmetal oxide => salt
Decomposition: A reaction in which a substance breaks down into to or simpler substances.
**Opposite of combination!
Single Replacement: One element replaces another element in a compound.
Reactivity sequence:
Can react with acids, not with H2O
Metals: Li>K>Ba>Ca>Na>Mg>Al>Zn>Fe>Ni>Sn>Pb>H>Cu>Hg>Ag>Au
can react with H2O and acids <=
=> cannot react with acids or H2O
Nonmetals: F>O>Cl>N>Br>S>I
Double Replacement: A reaction in which there is an exchange of components
AB+ CD → AD + CB
Complete Combustion of Organic Compounds: A reaction of organic compounds with oxygen. The result
is a production of water vapors, carbon dioxide gas, heat and flame.
Organic compound + Oxygen => carbon dioxide + water
Precipitate: A product of an aqueous reaction that settles to the bottom of the reaction container.
Acid: H+ + Acid remainder
Base: Metal Ion + OHSalt: Metal + Acid remainder
Oxides: Element + Oxygen
IDEAS FROM UNIT
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See compound naming flow chart
Completing and balancing all types of reactions.
Ionic equations with solubility rules.
SECTIONS TO REFER TO/ DOCUMENTS
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Chapter 2 (pg. 48-72)
Flow chart
Kinds of inorganic compounds sheet
Reaction types summary sheet.
Reaction Types Questions
1. What is K2O2?
2. What is the formula for Ammonium nitrate?
3. Setup for combination/composition reactions?
4. Basic Setup for single replacement reactions?
5. Basic setup for decomposition reactions?
6. Setup for double replacement reactions?
7. Setup for complete combustion reactions?
8. What are the products of complete combustion reactions?
9. What is a binary compound?
10. Complete the equation Al + H2SO4 
11. What is ionic bonding?
12. What is covalent bonding?
13. Name of Cr2O7-2?
14. What is a noble gas?
15. What is an oxide?
16. Balance:
NH3 + O2 
H2O +
NO2
17. What is the formula for Ammonia?
18. Write the formula for lithium oxalate?
19. What is the name of N2O?
20. What is the formula for Iron (III) chlorite?
21. Complete the decomposition equation for Al(OH)3
22.Balance the equation:
23. Balance the equation
PCl3 +
C+
24. What is the name Na2O?
H2O =>
Al2O3 =>
HCl +
Al4C3 +
H3PO3
CO
25. What is the name for MgCO3?
26. What is the formula for thiocyanic acid?
27. What is the charge of carbonate ion?
28. What is the name of H2CO3
29. Write in the formula and balance, hydrogen sulfide + ammonium chloride → hydrochloric acid +
ammonium sulfide?
30. What is the formula for perchloric acid?
31. What are the names of the following compounds?
CrSO4
FePO4
NH4Cl
32. What are the formulas for the following compounds:
periodic acid, sodium hydrogencarbonate, potassium hypoarsenite?
33. Which of the following is a cyanide? Mg3As2, NaCN, or KNO3
34. Match the following names with their ions:
C2O42- ___
(A) Thiosulfate ion
SCN- ___
(B) Thiocyanate ion
2S2O3 ___
(C) Dichromate ion
Cr2O72- ___
(D) Oxalate ion
Single Replacement:
Write the reaction. If no reaction, write NR
35. Mercury(II) chloride mixed with iodine
36. Lithium mixed with calcium oxide
37. Hydrogen mixed with copper(I) iodide
Double Replacement (All aqueous solutions)
Write the reaction. If no reaction, write NR
38. Sodium sulfate mixed with ammonium acetate
39. Potassium hydroxide mixed with carbonic acid
40. Magnesium chloride mixed with francium sulfide
41. Write the reaction of potassium oxide and sulfur trioxide
42. Write the reaction of vanadium (V) oxide and water
43. SX2 + H2O  H2SO3 What is X?
44. Write the reaction of the decomposition of sodium carbonate?
45. Write a balanced equation for the combustion of octane (C8H18) in oxygen.
Are the following compounds soluble in water?
46. ZrC2H3O2
47. La2(SiO4)3
48. Ra(OH)2
Balance the equations:
49. ____K2(Cr2O7) + ____HCl  ____Cl2 + ____H2O + ____CrCl3 + ____KCl
50. ____Cu + ____HNO3 ____H2O + ____NO + ____Cu(NO3)2
Answer Key (Naming, Reactions)
1. Potassium peroxide
2. NH4NO3
3. A + B
AB
4. AB + C
CB + A
5. AB A + B
6. AB + CD
AD + CB
7. CxHyOz + O2
CO2 + H2O
8. CO2 + H2O
9. A compound that consists of two elements.
10. 2Al + 3H2SO4
Al2(SO4)3 + 3H2
11. Attraction between positively and negatively charged ions.
12. A bond between two atoms which is formed by sharing of electrons between the atoms.
13. Dichromate ion
14. Elements in Group VIII and the elements which are least reactive.
15. Element + Oxygen
16. 4NH3 + 7O2
6H2O + 4NO2
17. NH3
18. Li2C2O4
19. Nitrogen oxide.
20. Fe(ClO2)3
21. 2Al(OH)3
Al2O3 + 3 H2O
22. PCl3 + 3H2O
3HCl + H3PO3
23. 9C + 2Al2O3
Al4C3 + 6CO
24. Sodium oxide
25. Magnesium carbonate
26. HSCN
27. (2-)
28. Carbonic Acid
29. H2S + 2NH4Cl
2HCl + (NH4)2S
30. HClO4
31. chromium (II) sulfate, iron (III) phosphate, ammonium chloride
32. HIO4, NaHCO3, K 3AsO2
33. NaCN, it is sodium cyanide
34. D, B, A, C
35. HgCl2 + I2
NR
36. 2Li + CaO
Li2O + Ca
37. H2 + 2CuI
2HI + 2Cu
38. Na2SO4 + 2NH4C2H3O2
2NaC2H3O2 + (NH4)2SO4 NR
39. 2KOH + H2CO3
K2CO3 + 2H2O
40. MgCl2 + Fr2S
2FrCl + MgS
41. K2O + SO3
K2SO4
42. V2O5 + H2O
V(OH)5
43. Oxygen
44. Na2CO3
Na2O + CO2
45. 2C8H18 + 25O2
16CO2 + 18H2O
46. Yes
47. No
48. Yes
49. K2(Cr2O7) + 14 HCl
3 Cl2 + 7 H2O + 2 CrCl3 + 2 KCl
50. 3 Cu + 8 HNO3
4 H2O + 2 NO + 3 Cu(NO3)2
Moles, Stoichiometry & Solutions
KEY TERMS
Actual yield: The amount of product actually obtained in a chemical reaction.
Anhydrous compound: A chemical compound that does not contain water molecules inside of crystals.
Aqueous solution: A solution in which water is the solvent.
Avogadro’s number: A number (6.022 x 1023) equal to the number of atoms in exactly 12.0 g of carbon-12;
the number of atoms, molecules, or formula units in one mole of an element.
Concentrated solution: A solution that contains a large amount of solute relative to the amount that could
dissolve.
Concentration: The measure of the quantity of a solute dissolved in a given quantity of solution.
Dilute Solution: A solution that contains a small amount of solute relative to the amount that could dissolve.
Empirical formula: A chemical formula that shows the lowest relative number of atoms of each element in a
compound.
Formula mass: The sum of the atomic masses (atomic weights in amu) of the atomic species as given in the
formula of the compound
Hydrate: A compound in which a specific number of water molecules are associated with each formula unit.
Hydrated ion: An ion surrounded with water molecules in a solution.
Hydration: Solvation in water.
Limiting reactant: The reactant that is consumed when a reaction occurs and therefore the one that
determines the maximum amount of product that can form.
Molarity (M): A concentration term expressed as the moles of a solute dissolved in 1L of solution.
Molar mass (g/mol): The mass of 1 mol of entities of a substance.
Mole: The SI unit based on the amount of a substance. One mole is the amount of substance that contains a
number of entities equal to the number of atoms in 12.0 g of carbon-12.
Molecular formula: A formula that shows the actual number of atoms of each element in a molecule
Molecular weight: The sum (in amu) of the atomic masses of a formula unit of a compound.
Percent by mass: The fraction by mass expressed as a percentage. A concentration term expressed as the
mass in grams of a solute dissolved per 100. g of solution.
Percent by volume: A concentration term defined as the volume (L) of a solute in 100.L of solution.
Percent yield: The actual yield of a reaction expressed as a percent of the theoretical yield.
Saturated solution: A solution that contains the maximum amount of dissolved solute at a given temperature
in the presence of undissolved solute.
Solute: The substance that dissolves in the solvent.
Solvent: The substance in which the solute dissolves.
Stoichiometry: The study of the mass-mole-number relationships in chemical formulas and reactions.
Theoretical yield: The amount of product predicted by the stoichiometrically equivalent molar ration in the
balanced equation.
Weak electrolyte: A weak electrolyte is a substance that dissociates into ions only to a small degree.
IDEAS FROM UNIT
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Quantity of Matter : The mole
Molecular mass and molar mass
Percent composition of compounds
Empirical formula
Quantitative meaning of a chemical equation
Calculation of the amounts of reactants and products
Limiting Reagent
Percent Yield
Concentration of solutions
Dilution of solutions
Concentrations of ions in solutions
SECTIONS TO REFER TO/ DOCUMENTS
 Chapter 3 (pg. 87-133)
 Stoichiometry In-Class Problems
 Stoichiometry Worksheet
 Limiting Reagent Notes
 Solution Concentration Worksheet
Stoichiometry Questions
1. What is the atomic weight of an element? How does the atomic weight of an element differ from the
mass number of an isotope?
2. How are atoms weights determined?
3. What is the SI unit for quantity of matter?
4. What is Avogadro’s number?
5. What is molecular weight? Formula weight? Molar mass?
6. What advantage is there in using a mole as a unit for quantity of matter rather than a kg or pound?
7. What is a hydrate?
8. List some quantitative ways to express the concentration of a solution.
9. What is a limiting reactant in a reaction?
10. Why is it called the limiting reactant?
11. What additional information other than the percent composition would you need to determine the
molecular formula of a compound?
12. How many atoms does a mole of atoms contain?
13. How many molecules are in a mole of molecules?
14. Calculate the percent composition of the elements in the following compound, CuSO4(H2O)5
15. A 1.00-gram sample of a compound loses .450 grams of oxygen when it decomposes upon heating.
The solid residue that remains is sodium chloride. Derive the simple formula for the compound.
16. When a 1.214-gram sample of a compound consisting of only carbon and hydrogen is burned, 4.059
grams of carbon dioxide and .9494 grams of water are produced. Calculate the percent of carbon and
percent of hydrogen in the compound and derive the empirical formula for the compound.
17. H2(g) reacts with O2(g) to form water vapor according to the equation: 2H2(g) + O2(g)
there are 5.60 g of H2 how many grams of O2 and H2O are there?
2H2O(g). If
18. Copper metal reacts in concentrated sulfuric form hydrated copper sulfate, CuSO4(H2O)5,
according to the equation Cu(s) + 2H2SO4(aq) + 3H2O(l)
CuSO4.5H2O(s) + SO2(g). How many grams of
hydrated copper sulfate can be obtained in this reaction when 20.0 g of 98.0 percent sulfuric acid reacts
with excess copper, assuming that the yield is 85.0 percent?
Given the following equation: 2 KClO3
2 KCl + 3O2
19. How many moles of O2 can be produced by letting 80. grams of KClO3 react?
20. 4.64L CO2 at STP reacts with excess CaO to produce 24.5 g of product, what is the percent yield?
21. Calculate the empirical formula for a compound that has 43.7 g phosphorus and 56.3 grams of oxygen.
22. Hydrogen gas and nitrogen gas are reacted to produce NH3 gas. If you have 2.34 g of H2 how many grams
of NH3 can you produce? (Hint: write out equation!)
Given the equation 2Al + 6HCl = 2AlCl3 + 3H2
23. What is the Limiting reagent if 10.0g of Al and 5.0g of HCl react?
24. How many moles of AlCl3 do you have?
25. 0.0200 mol of a gas occupy 2.00L at certain conditions. 200.L of another gas, at the same conditions,
weighs 8.00 g. What is the molar mass of the second gas? What is the element?
Use this equation: 2HCl + Na2CO3
2NaCl + H2CO3
26. If 32g HCl react with 20g of Na2CO3, what is the limiting reagent? How much H2CO3 is produced, in
grams?
For the equation C4H10O7 + 3O2
4CO2 + 5H2O
27. What mass of CO2 is produced when 36.0 g C4H10O7 is burned?
28. For the same reaction, what mass of H2O is produced when 10L of O2 is burned at STP?
For the reaction CaO + CO2
CaCO3
29. How much CaCO3 is made, in grams, when 12.0 g CO2 reacts with excess CaO?
30. Iron reacts with chlorine gas to form iron chloride. If 30.0 g Fe react with 50.0 g Cl2, what is the limiting
reactant? How much FeCl3 is produced?
31. What is the molarity of a sodium chloride solution that contains 12.4 g NaCl in 350. mL of solution?
32. Find the mass percent of each element in sucrose (C12H22O11)
In 1.28 mol Ag...
33. …how many grams of Ag?
34. …how many Ag atoms?
35. Write a balanced equation for the complete combustion of glucose (C6H12O6)... if 72.0 g glucose
combusts with excess O2, what is mass of CO2 formed?
More Stoichiometry questions
36. What is the percent by mass of oxygen in Fe2O3?
37. The human body needs at least 1.03 x 10-2 mol O2 every minute. If all of this oxygen is used for the
cellular respiration reaction that breaks down glucose, how many grams of glucose does the human body
consume each minute?
C6H12O6(s) + 6 O2(g) -----> 6 CO2(g) + 6 H2O(l)
38. Carbon monoxide can be combined with hydrogen to produce methanol, CH3OH. Methanol is
used as an industrial solvent, as a reactant in synthesis, and as a clean-burning fuel for some racing
cars. If you had 152.5 kg CO and 24.50 kg H2, how many kilograms of CH3OH could be produced?
39. Dimethylglyoxime is a molecular compound with many uses. An elemental analysis of
Dimethylglyoxime shows is to be 41.37% carbon, 6.94% hydrogen, 24.12% nitrogen, and 27.56% oxygen by
weight. If Dimethylglyoxime contains only these elements, what is its empirical formula?
40. In the space shuttle, the CO2 that the crew exhales is removed from the air by a reaction within canisters
of lithium hydroxide, which creates water that is consumed by the astronauts. On average, each astronaut
exhales about 20.0 mol of CO2 daily. If there are five astronauts and there is only 500 Kg of Lithium
Hydroxide, how long will it be before their supply of water stops growing?
CO2(g) + 2 LiOH(s) ------> Li2CO3(aq) + H2O
Answer Key (Moles & Stoichiometry)
1. The average of the masses of its naturally occurring isotopes weighted according to their abundances.
2. The natural abundance of each isotope of an element is considered to determine the weights.
3. A mole
4. 6.022 x 1023
5. - The molecular weight is the sum of the atomic weights of the atoms in a molecule of a compound.
- The formula weight is the sum of the atomic weights of the atoms in a formula unit.
- The molecular mass is the mass of one mole of any substance.
6. The advantage of using moles is that the quantity will be much smaller and a number that is easier to deal
with than if you use grams or pounds. Also, you can compare two quantities of moles to each other, but you
cannot compare grams and pounds.
7. Hydrates are compounds formed by the union of water with some other substance, generally forming a
neutral body, as certain crystallized salts.
8. The concentration of a solution is usually given in moles per liter (mol x L-1 OR mol/L). This is also known
as molarity.
9. In chemistry, the limiting reagent is the chemical that determines how far the reaction will go before the
chemical in question gets used up, causing the reaction to stop.
10. It is called the limiting reagent because it limits the amount of product that can be produced
during a reaction.
What is… (give an example of each)
11. To obtain the molecular formula for a compound, its molecular weight is compared with the weight of the
empirical formula. The simplest formula of the compound can be obtained from the mole ration of the
elements, which are the same as their atomic weight.
12. Avogadro’s number: 6.022 x 1023
13. Avogadro’s number: 6.022 x 1023
14. 1 mol Cu x 63.55 g of Cu/ 1 mol = 63.55 g Cu
1 mol S x 32.06 g of S/ 1 mol = 32.06 g S
9 mol O x 16 g of O/ 1 mol = 144 g of O
32.06/249.69 = 13%
144/249.69 = 58%
10 mol H x 1.008 g of H/ 1 mol = 10.08 g of H
15. NaxClzOy
63.55/249.69 = 25%
10.08/249.69 = 4%
NaCl
1-.450 = .55g of NaCl .45gO x 1 mol O /16 g O = .02813 mol O
.55 g of NaCl x 1mol/58.44 g NaCl x 1molCl/1molNaCl x 35.45gCl/1 molCl =.3337 g of Cl = .009411 mol Cl
.55g -.337g = .2163 g Na .2163gNa x 1 mol/22.98g Na = .009412 mol Na
NaClO3
16. CxHy + O2
CO2 + H2O
4.055 g CO2 x 1 mol CO2/44.01 g CO2 x 1mol C/1 mol CO2 x 12.01 g C/1 mol C = 1.107 g C = .09223 mol C
.9494 g H2O x 1mol H2O/18 g H2Ox 2 mol H/1 mol H20 x 1.008 g H/ 1 mol H = 1.063 g H = .1055 mol H
0.09223mol C:0.1055mol H = 1:1
1.108g C + 0.1063g H=1.214g
CH 91.26% C, 8.756% H
CH
1.108g C/1.214g=91.26% C
0.1063g H/1.214g=8.756% H
17. 5.6 g H2 x 1 ml H2/2.02 g H2 x 1 mo O2/2 mol H2 x 32 g O2/1 mol O2 =44.36 g O2
5.6 g H2 x 1 mol H2/2.02 g H2 x 1 mol of H2O/1 mol of H2 x 18 g of H2O/1 mol H2O = 49.90 g H2O
18. 20gH2SO4 / 98.08gH2SO4 x 1 mol CuSO4(H2O)5 / 2 mol H2SO4 x 249.61 g CuSO4(H2O)5= 25.45 g
CuSO4(H2O)5
25.45 g CuSO4(H2O)5 x .85 = 21.63 CuSO4(H2O)5
Given the following equation: 2 KClO3
2 KCl + 3 O2
19. (.97 mol O2)
20. CO2 + CaO
CaCO3
(20.73 g CaCO3)
(Percent yield = 84.6%)
21. 3.52/1.41=2.5 (THE EMPIRICAL FORMULA IS:P2O5)
22. (13.2 grams NH3)
Given the equation 2Al + 6HCl = 2AlCl3 + 3H2
23. (LR= HCl. )
24. (mol AlCl3 =0.046)
25. 4.00g/mol Helium
Use this equation: 2HCl + Na2CO3
2NaCl + H2CO3
26. 22.1 g NaCl Na2CO3 is LR
For the equation C4H10O7 + 3O2
4CO2 + 5H2O
27. 37.3 g CO2
28. 13.4 g H2O
29. 27.3 g CaCO3
30. Fe is limiting reactant, 87.0g FeCl3 formed.
31. 0.606 M
32. 42.1%C, 6.48% H, 51.4%O
In 1.28 mol Ag...
33. 138 g Ag
34. 7.71 X 1023 Ag atoms
35. C6H12O6 + 6O2
6CO2 + 6H2O
106 g CO2
38. 174 Kg (CH3OH)
39. C2H4NO
36. 30.0% (O)
40. 67.7 days
37. 31.2 g (C6H12O6)
Energy
KEY TERMS
Endothermic reaction: A reaction that absorbs heat from its surroundings and therefore increases the
enthalpy of the system.
Exothermic reaction: A reaction that emits heat into its surroundings and therefore decreases the enthalpy of
the system.
Calorimetry: A method of determining the energy exchange between the reaction system and its
environment.
Enthalpy (H): The thermodynamic quantity that the sum of the internal energy plus the product of the
pressure and volume.
Specific heat capacity (c): The quantity of heat is required to change 1 gram of a substance by 1ºC.
Temperature: The measure of average molecular kinetic energy.
IDEAS FROM UNIT
 ΔH
 Enthalpies of formation/reaction
 Hess’s Law of heat summation
 Q = mcΔT
 4.184 J = 1 cal
SECTIONS TO REFER TO/ DOCUMENTS
 Chapter 6 (pg. 225-255)
 Energy Worksheet
Energy Questions
1. What is the meaning if a positive sign for the enthalpy of a reaction?
2. Define heat capacity.
3. Define specific heat.
4. A reaction absorbs 15.0 Joules of heat. How many calories is this?
5. If the heat of a reaction is -20 Joules, is this reaction endothermic or exothermic?
6. Define a calorie in terms of the quantity of water in its temperature change:
7. What is ΔH?
8. Write the formula of ΔH.
9. What is the first law of thermodynamics?
10. What is Hess’s Law?
11. How is ΔHrxn related to enthalpies of formation of the reactants and the products?
12. It took 478 J to heat 199.9 g of "Z" from 34.62ºC to 46.64ºC. What was the specific heat of "Z"?
13. How many kJ of energy will be released when 12.50 L of CH4 (gas at SATP) is burned?
1CH4 + 2O2
1CO2 + 2H2O + 890. kJ
14. 1S + 1O2
1SO2
1 SO2 + 1/2 O2
1SO3
1 SO3 + 1 H2O
1 H2SO4
∆H = -297kJ/mol
∆H = -144kJ/mol
∆H = -84kJ/mol
Calculate the ∆H for the reaction 1 S + 3/2 O2 + 1 H2O
1 H2SO4
15. Use the following enthalpies of formation to calculate the ∆H for the following reaction and indicate
whether the reaction is exo- or endothermic.
3 C + 1 SiO2  1 SiC + 2 CO
SiC: -27 kJ/mol, CO: -26 kJ/mol, SiO2: -205 kJ/mol
16. How many grams of C2 H4 must be burned in oxygen to release 1550 kJ? (Equation? ∆H?)
17. How many joules of heat is released when 2.0 L of water is cooled from 40ºC to 15ºC?
18. The specific heat capacity of copper is 0.38 J g -1 ºC-1. What is the temperature increase of 2.0 kg of copper
after absorbing 8.4 kJ of heat?
19. 4 Fe(s) + 3 O2(g)
2Fe2O3 (s)
∆Hrxn = -1.65 x 10³ kJ/mol
How much heat is involved when 5.00kg of iron reacts?
Answers (Energy):
1. This means energy was absorbed during the reaction.
2. The amount of heat required to raise the temperature of a gram of a substance by 1 degree Celsius.
3. The heat capacity of 1 gram of a substance.
4. 3.59 calories
5. Exothermic
6. One calorie is the heat/energy required to raise 1g H2O 1 degree Celsius.
7. The sum of the internal energy change of a system and its pressure x volume is called delta H enthalpy
change.
8. ∆H= ∆E + P ∆V at constant pressure
9. Energy is neither created nor destroyed.
10. Enthalpy changes of different steps of a reaction can be added to obtain the delta H for the reaction.
11. ∆Hrxn = ∆Hf (products)- ∆Hf (reactants)
12. C= Q/M∆T
C= 478 J / (199.9 g x 12.02 g) = .199 J/g x degrees Celsius
.199 J/g x degrees Celsius
13. 12.50 L CH4 x 1 mol CH4 x 890. kJ = 449 kJ
24.8 L CH4 1 mol CH4
14.
1 S + 1 O2
1SO2
∆H = -297Kj/mol
1 SO2 + 1/2 O2
1SO3
∆H = -144Kj/mol
1 SO3 + 1 H2O
1 H2SO4 ∆H = -84Kj/mol
-297 + -144 + -84 = -525
∆H = -525 kJ/mol S (*remember: use whatever product/reactant there is one whole mole of (i.e. 1S as
opposed to 1/2 O2 when giving the unit of mol/x.)
15. Reaction:
3 C +1 SiO2
1 SiC + 2CO
0 + -205
-27 + -52 = +126 KJ/mol SiO2
+126 KJ/mol SiO2 ; Endothermic reaction
16. 1550 kJ x 1 mol C2 H4 x 28.4 C2 H4
1412 KJ
1 mol C2 H4 = 30.7 g C2H4
30.7 g C2H4
17. 2.1 x 105 joules
18. 11.1 degrees Celsius
19. 5.00kg Fe x 1000g Fe x 1mol Fe x 2mol Fe2O3 x -1.65 x 10³ kJ =
1
1kg Fe
55.85g Fe 4mol Fe
1mol Fe2O3
(3 significant figures): 73,900kJ
73,858.55 kJ
Gas Laws
KEY TERMS
Pressure (P): the force exerted per unit of surface area. P= force/area.
1 atm = 101.325 kPa = 760 mmHg = 760 torr
Barometer: common device used to measure atmospheric pressure.
Manometer: device used to measure the pressure of a gas in an experiment.
Pascal (Pa): SI unit of pressure.
Standard atmosphere (atm): another unit of pressure (large).
Millimeters of mercury (mmHg): common pressure unit based on measurement with a barometer or
manometer.
Ideal gas: a theoretical gas that has no shape or intermolecular forces involved, exhibits simple linear
relationships among volume, pressure temperature and amount
(n= # of moles, P= pressure, T= temperature [in K], V= volume [L/mL])
Daltons’s Law: n vs. P, direct proportion
Gay-Lussac’s Law: P vs. T, direct proportion
Charles’s Law: V vs. T, direct proportion
Unnamed Law: n vs. T, indirect proportion
Boyle’s Law: P vs. V, indirect proportion
Avagadro’s Law: n vs. V, direct proportion
Two-Condition Equation:
P 1 V1
n1T1
= P2V2
n2T2
From the above relations, we derive:
Ideal Gas Law: PV = nRT
Universal gas constant (R): a proportionality constant
8.314 kPa•L/mol•K= 0.08206 L•atm/mol•K= 62.36mmHg•L/mol•K
Standard temperature and pressure (STP): 0°C (273.15K) and 1 atm (760 mmHg)
Standard molar volume: 22.4L per mole of any gas
Partial pressure: portion of the total pressure of the mixture that is the same pressure it would exert by itself.
Dalton’s Law of partial pressure: in a mixture of unreacting gases, the total pressure is the sum of the
partial pressures of the individual gases
Mole fraction: component in a mixture contributes a fraction of the total number of moles in the mixture
Density (D): mass per volume of a substance
Molar Mass Equation: a derivation from the ideal gas law, we get:
Molar Mass = DRT
P
Kinetic molecular theory: draws conclusions through mathematical derivations
RMS speed: A molecule moving at this speed has the average kinetic energy
Effusion: the process by which a gas escapes from its container through a tiny hole into an evacuated
Graham’s Law of effusion: the rate of effusion of a gas is inversely proportional to the square root of its
molar mass:
Diffusion: the movement of one gas though another
SECTIONS TO REFER TO/DOCUMENTS
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Chapter 5 (pg. 177-223)
Gas worksheets 1-5
“It’s a Gas” sheet
Gas Law Questions
1. What is the density of a gas that is 1.996g with a volume of 1190cm3?
2. 3.200 g of a gas is 5.080 mol. What is its molar mass?
3. Oxygen gas effuses through a capillary in 2.00 seconds. An unknown gas of the same volume effuses in
4.00 seconds. What is the molar mass of the unknown gas?
4. In a closed end manometer, the mercury level was 540.mm higher on the closed end than on the gas end.
What was the gas pressure?
5. If the Thirsty Bird (think Gay Lussac’s Law) dips his head in water and cools down from 50.0◦C to 30.0◦C,
and the pressure in his belly was originally 1.00atm, what was the pressure after his drink?
6. A snake balloon has .00100 moles of air in it and a volume of .100mL. If I blow .028 moles of air into it,
what is its volume?
7. A He molecule, weighing 4.0 g/mol at 200K has ____times as much K.E. as an H2 molecule, weighing 2.0
g
/mol, at 200K?
8. H2 molecules (2.00 g/mol) at 200.K go ____as fast as He molecules (4.00 g/mol) at 600.K?
9. Charles’ law: a gas with a volume of 700.mL at 300.K has what volume at 383K?
10. In an open end manometer with an atmospheric pressure of 37.8kPa, the Hg level was 27.3 mm higher on
the left. What is the gas pressure?
11. At 27 degrees Celsius, a gas, was at 321kPa, at what temperature would it be at 1.00atm?
12. 7.36 grams of a gas occupies 2.67 liters at 17 degrees Celsius and 115.2kPa. How many moles of gas is it?
13. What is the density of CO2 at 731 K in 7.32 psi?
14. What is the molecular weight of a gas that has a mass of 13.2 grams, a volume of 3.7 liters, at 32 degrees
Celsius, and 117kPa?
15. 113.cm3 of gas measured at 183.kPa, 343.K, would occupy what volume at SATP?
16. What is the volume of 0.936 moles of gas at 347K, 117kPa?
17. What volume of oxygen gas, measure at 35OC and 115kPa is required to completely burn 50.0g of Mg?
18. What is the volume of 14.7 g of F2 gas at -15.8OC, 187kPa
19. At what temperature will 1.43 moles of gas produce a pressure of 132.kPa in a 440.cm3 flask?
20. A molecule at 12.0K has ____ as many times kinetic energy as a molecule at 4.00K.
21. Cl2 molecules at 243 K go _____ X as fast as O2 molecules at 513K.
22. What is the molecular weight of a gas that has a density of 4.01g/L at 15°C and 876 mmHg?
23. What is the density of SO2(g) at 923K, 675mmHg?
24. What volume does 65L of a gas at 24atm occupy at 54atm?
25. At 253K a gas had a pressure of 1.36atm. What temperature would it be at 3.46atm?
26. 3.0 moles take up 46 L. How many moles would occupy 26L at the same temperature and pressure?
27. 0.336 moles of gas produced 256kPa. What pressure would result from 0.662 mol?
28. 0.567 mol of gas in a container at 321.1K produces certain pressure. What temperature would be needed
to produce the same pressure with 0.223 mol or gas in the same container?
29. In a closed end manometer, the mercury level was 673 mmHg higher on the closed end than the gas side.
What was the pressure of the gas?
30. In an open end manometer, atmospheric pressure is 756mmHg and the Hg level is 83 mm higher on the
left. What is the pressure of the gas?
31. In an open end manometer with atmospheric pressure at 96.5kPa, the mercury level is 46mm higher on the
right. What is the pressure of the gas?
32. 0.575 mol produced 1.00atm at -30°C. How many moles of gas must be released from the container in
order to keep the pressure the same when the temperature is increased to 67°C?
33. What is the density of CH4(g) at STP?
34. What is the density of SO2(g) at 822K, 0.281atm?
35. If a gas with a density of 0.264 g/L is at 729 mmHg at a certain temperature and number of moles, what
would a gas with a density of 0.645 g/L be at for the same condition?
Answers (Gas Laws)
What is the density of a gas that is 1.996g with a volume of 1190cm3? 1.677 x 10-3 g/cm3
2. 3.200 g of a gas is 5.080 moles. What is its molar mass? 6.299 x 10-1 g/mol
3. Oxygen gas effuses through a capillary in 2.00 seconds. An unknown gas of the same volume effuses
in 4.00 seconds. What is the molar mass of the unknown gas? 128 g/mol
4. In a closed end manometer, the mercury level was 540.mm higher on the closed end than on the gas
end. What was the gas pressure? 540. mmHg
5. If the Thirsty Bird dips his head in water and cools down from 50.0◦C to 30.0◦C, and the pressure in
his belly was originally 1.00atm, what was the pressure after his drink? 0.938atm
6. A snake balloon has .00100 moles of air in it and a volume of .100mL. If I blow .028 moles of air into
it, what is its volume? 2.9mL
7. An He molecule, weighing 4.0 g/mol at 200K has ____times as much K.E. as an H2 molecule,
weighing 2.0 g/mol, at 200K? 1 (the same amount of K.E.)
8. H2 molecules (2.00 g/mol) at 200.K go ____times as fast as He molecules (4.00 g/mol) at 600.K? 3
9. Charles’ law: a gas with a volume of 700.mL at 300.K has what volume at 383K? dd. 0.834 L
11. In an open end manometer with an atmospheric pressure of 37.8kPa, the Hg level was 27.3 mm higher
on the left. What is the gas pressure? 257 mmHg
12. At 27 degrees Celsius, a gas, was at 321kPa, at what temperature would it be at 1.00atm?
13. 7.36 grams of a gas occupies 2.67 liters at 17 degrees Celsius and 115.2kPa. How many moles of gas
is it?
14. What is the density of MgO at 731 K in 7.32 psi?
15. What is the molecular weight of a gas that has a mass of 13.2 grams, a volume of 3.7 liters, at 32
degrees Celsius, and 117kPa?
16. 113.cm3 of gas measured at 183.kPa, 343.K, would occupy what volume at SATP? V2 = 174.cm
17. What is the Volume of 0.936 moles of gas at 347K, 117kPa? V=22.4 L
18. What volume of oxygen gas, measured at 35OC and 115kPa is required to completely burn 50.0g of
Mg? 22.9 L O
19. What is the volume of 14.7 g of F2 gas at -15.8OC, 187kPa? V= 3.90L
20. At what temperature will 1.43 moles of gas produce a pressure of 132.kPa in a 440.cm3 flask?
T=4.89K
21. 3.00
22. 0.462
23. 8.22 g/mol
24. 0.751 g/L
25. 29 L
26. 644 K
27. 1.70 mol
28. 816K
29. 1480K
30. 673 mmHg
31. 673 mmHg
32. 770. mmHg
1.
Convert to Kelvin:
33. 248K
34. 420 K
35. 241.2K
36. 296 K
37. 254 K
38. 38. 0.542atm
39. 5720 mmHg
40. 1420kPa
41. 0.164 mol
42. 0.645 g/L
43. 0.267 g/L
44. 298 mmHg
Atomic Structure and Periodicity
KEY TERMS
Electronegativity: the relative ability of an atom in a molecule to attract bonding electron pairs
Periodic trend: how a certain aspect of periodicity changes as you move down and across the periodic table
Ionization energy: the amount of energy needed to remove an outer electron from its orbit periodic trend:
increases to the right, decreases going down
Electron affinity: energy change that occurs when an electron adds to an isolated atom to form a negative ion
periodic trend: increases left to right, decreases going down
Cathode: A negatively charged electrode
Anode: A positively charged electrode
Frequency: The number of waves that pass a given reference point in a unit time
Hertz (Hz): The SI unit for frequency which equals one wave per second, therefore 1 Hz = 1 wave/s
Quantum levels: Energy levels
Ground state orbit: Orbit closest to the nucleus
Principal quantum levels: shells
Plank's Constant (h): 6.626 x 10-34, used to find frequency
Valence electrons: The electrons of an atom that can be involved in chemical bonding; the electrons on the
outermost energy level.
Atomic radii: The radius of an atom which follows this periodic trend: the radius decreases left to right
within a period, increases going down.
Formulas:
c = λ·ν
speed of light=wavelength·frequency
c = speed of light = 3.00x10^8 m/s
λ = wavelength
ν = frequency
∆E = ∆nhν
∆E = change in energy
∆n = change in quantum number
h = Planck’s constant=6.626x10^-34 J·s
ν = frequency
λ=h/mu
wavelength = Planck’s constant/(mass·speed)
λ=wavelength
h=Planck’s constant
IDEAS FROM UNIT
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Electron Configuration
Orbital Box Diagrams
Lewis (e-) Dot structures
Probability Diagram
Valence Electrons
Quantum Numbers
Periodic Trends
Waves and Frequency
Bohr Model
Calculation of energy released from electrons jump from one shell to another
SECTIONS TO REFER TO/ DOCUMENTS
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Chapter 7 and 8 (pg. 257-327)
Lewis Dot Structure Rules
Types of Substances
Atomic Structure and Periodicity Questions
1. What is the formula for frequency?
2. What is the formula for energy?
3. What kind of wave carries the most energy?
4. What is the visible light spectrum?
5. What are the levels called that electrons get excited to?
6. What are the known four subshell letters?
7. What are the quantum levels?
8. How many electrons can fit in a p subshell?
9. What are the two spin quantum number values?
10. How many electrons can fit in an f subshell?
11. How would you classify elements into groups with a whole bunch of similar properties?
12. How would you classify elements with practically the same properties?
13. How would you classify elements by the order of filling electron orbitals?
14. How would you classify elements by the electron sublevels being filled?
15. What is the periodic trend of ionization energy? Why?
16. What is the difference between 1st and 2nd ionization energies?
17. Rank from lowest to highest ionization energy: F, Se, Ra, Pd, Ne.
18. Which is greater, 1st or 2nd ionization energy?
19. What is the trend in the periodic table for atomic size?
20. Rank from smallest to largest atomic size: Ni, Ge, Pd, Y, P, F, Ba, Ra, Cl.
21. Rank in order of increasing electron affinity: Mn, C, He, B, La, W, Ni
22. Which has the higher electron affinity, Co or At?
23. True or False: The periodic trend of atomic size is that it decreases as you go down because the number
of energy levels decreases. __________
Fill in the blanks:
24. _________ ions are always smaller than the parent atoms and ________ ions are always larger.
25. Noble gases have a very_______ ionization energy.
26. Which takes precedence, period number or group number?
27. An atom of which element will have a highest energy electron with the following set of quantum
numbers? (4. 1. -1. -1/2.)
28. Write the Electron configuration for a Cu atom using abbreviation.
29. Which elements have nothing but inner core electrons?
30. A main-group element with a (MS +1/2) quantum number would be located in
a. group 2, 3, 4 or 5 b. group 1, 6, 7 or 8 c. group 2, 6, 7 or 8 d. group 1, 3, 4 or 5
31. What are the quantum numbers of the outer most electron of the most reactive metal?
32. How many electrons can fit on a d subshell?
a. 10 b. 4 c. 6
d. 8
33. How many orbitals are there on a p subshell?
a. 1
b. 3 c. 4
d. 5
34. How many electrons are there on an orbital?
a. 1
b. 2 c. 3
d. 4
35. How many electrons are there in one Li molecule?
a. 1
b. 2 c. 3
d. 4
36. What shape is an “s” subshell?
a. square
b. peanut
c. dodecahedron
d. sphere
37. What ion with a 2+ charge has the electron configuration 1s22s22p63s23p6?
38. Order the following (N, F, Rb, Cs) from smallest to largest atomic size:
from least to greatest electron affinity:
39. Which element, when oxidized and dissolved in water, turns blue litmus red, K or Br? Why?
40. Which of the following tend to produce 2+ ions? Na, Ne, Mg, S, Zn, Y
41. Which has a greater atomic radius, O or O2-? Why?
42. Give the quantum numbers for a 4d4 electron.
43. What is the electron configuration of Chromium?
A. 1s21p62s22p63s23p64s23d4
B. 1s22s22p63s23p64s13d5
C. 1s22s22p63s23p63d44s2
D. 1s22s22p63s23p63d54s1
44. If a wave has a frequency of 400.Hz, what is its’ wavelength in meters?
45. Match the electron configurations with the elements (not all choices will be used):
1. 1s22s22p4
____A. Uranium
2. 1s22s22p63s23p64s13d10
____B. Phosphorus
3. 1s22s22p63s23p1
____C. Copper
4. 1s22s22p63s23p64s23d104p6
____D. Zinc
2 2 6 1
5. 1s 2s 2p 3s
____E. Krypton
6. 1s22s22p63s23p3
____F. Oxygen
46. Assume it takes 1.64x10^-18 J for an electron to jump from energy level (n) 1 to 2, 1.94x10^-18 J from 1
to 3, and 2.04x10^-18 J from 1 to 4.
a. If an electron receives 1.93x10^-18 J, what energy level will it jump to?
b. If the electron jumps back to the ground state, what is the wavelength of the photon of light emitted?
47. Draw the lewis dot diagram for Zinc.
Answers (At. Structure & Periodicity)
1.speed of light divided by wavelength
2. hv
3. UV
4. 400-700nm
5. quantum levels
6. s, p, d, f
7. K, L, M, N
8. 6
9. +1/2, -1/2
10. 14
11. Metals, non-metals, metalloids
12. Families (alkali metals, alkaline earth metals, halogens, noble gasses)
13. Representative elements, transition metals, inner transition metals
14. s-elements, p-elements, d-elements
15. It decreases going down because added energy levels shield the power of the nuclear charge making it
easier to remove most electrons. It increases going to the right because the increase in nuclear energy makes it
more difficult to remove the outermost electrons.
16. 1st ionization energy is the energy needed to remove one electron. 2nd ionization energy is the energy
needed to remove a second electron.
17. Ra, Pd, Se, F Ne.
18. 2nd ionization energy
19. Smaller up and right. Larger down and left.
20. F, Cl, P, Ge, Ni, Pd, Y, Ba, Ra.
21. La, W, Mn, Ni, B, C, He.
22. Co
23. F
24. positive ions are always smaller than the parent atoms and negative ions are always larger.
25. period number
26. high
27. Selenium
28. [Ar] 4s1 3d10
To complete its D shell it steals an electron from the S orbital
29. Noble gases
30.C) group 2, 6, 7 or 8
31. Francium (7. 0. 0. +1/2.)
32. a. 10
33. b. 3
34. b. 2
35. c. 3
36. d. circle
37. Ca
38. size: F, N, Rb, Cs
electron affinity: Cs, Rb, N, F
39. Br, because acids turn blue litmus red, and nonmetals like Br tend to produce acids.
40. Mg, Zn
41. O2- will have a greater atomic radius, because the two extra electrons in O2- occupy outermost shells and
increase electron repulsion, expanding the radius.
42. n=4, l=2, m=+1, s=-1/2 (or +1/2)
43. b. 1s22s22p63s23p64s13d5
44. 7.50x10^5 m
45. B-6, C-2, E-4, F-1
46. a. n=2
b. 2.424x10^-7 meters
47. :Zn
Bonding
KEY TERMS
Single bond: a bond that consists of one electrons pair.
Double bond: a covalent bond that consists of two bonding pairs: two atoms sharing four electrons in the
form of one σ and one π bond.
Triple bond: a covalent bond that consists of three bonding pairs: two atoms sharing six electrons in the form
of one σ and two π bond.
σ (sigma) bond: a type of covalent bond that arises through end-to-end orbital overlap and has the most of its
electron density along the bond axis.
π (pi) bond: a covalent bond formed by sidewats overlap of two atomic orbitals that has two regions of
electron density, one above and one below the internuclear axis.
cis isomer: an isomer in which two atoms or groups of atoms are on the same side of some reference line or
plane in the molecule.
trans isomer: an isomer in which two atoms or two groups of atoms are on the opposite sides of some
reference line or plane in the molecule.
Lewis structure: A structure consisting of electron-dot symbol, with lines as bonding pairs and dots as lone
pairs.
Octet rule: The observation that when atoms bond, they often lose, gain, or share electrons to attain a filled
outer shell of eight electrons.
Molecular shape: The three-dimensional structure defined by the relative positions of the atomic nuclei in a
molecule.
Tetrahedral: a molecular shape formed when four electron groups maximize their separation around a central
atom.
See-saw: a molecular shape caused by the presence of one equatorial lone pair in a trigonal bipyriamidal
arrangement.
Flat triangle: a molecular geometry with one atom at the center and three atoms at the corners of a triangle
all in one plane.
T-shape: A molecular shape caused by the presence of two equatorial lone pairs in a trigonal bipyramidal
arraignment.
Isomers: Compounds with the same molecular formula but different arraignments of atoms.
Resonance form: One or two Lewis-dot structures that can be drawn for a molecule (with double or triple
bonds).
Bond angle: The angle between two lines drawn from the central atom of a molecule to two adjacent atoms
bonded to the central atom.
Dipole moment (μ): The product of partial charges on two atoms and the distance between the nuclei.
Valence shell: the outermost shell of an atom, which contains the electrons most likely to account for the
nature of any reactions involving the atom and of the bonding interactions it has with other atoms
Hybrid orbitals: An atomic orbital postulated to form during bonding by mathematical mixing of specific
combinations of non-equivalent orbitals in a given atom.
Van der Waals forces: Intermolecular forces that include both dipole-dipole interaction and London forces.
Hydrogen bonds: A type of dipole-dipole attraction that arises when between molecules that have an H atom
bonded to a small and highly electronegative atom with lone pairs (usually F, O, N).
VSEPR theory: A model explaining that the shapes of the molecules and ions result from minimizing
electron-pair repulsions around the central atom.
Lone pair: An electron pair that is part of an atom’s valence shell but not involved in covalent bonding.
Bond energy: The enthalpy change accompanying the breakage of a given bond in a mole of gaseous
molecules.
IDEAS FROM UNIT
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chemical intermolecular/intramolecular bonding
ionic bonding
covalent bonding
-non-polar
-polar
metallic bonding
orbital hybridization
orbital overlap
orbital box diagrams
VSEPR theory
Lewis dot structures
valence shells
bond polarity
molecular polarity
molecular shapes
isotopes
atomic properties
bond energy
chemical change
Hydrogen bonds
Van der Waal’s forces
SECTIONS TO REFER TO/DOCUMENTS
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Chapters 9-11 (pg. 329-488)
VSEPR chart
Types of Substances
Bonding Questions:
1. What is the hybridization of a central atom of SeCl6?
2. What kinds of orbitals form sigma and pi bonds?
3. Which is stronger, a sigma or a pi bond, and why is that so?
4. Name exceptions to the octet rule.
5. How many sigma and pi bonds does a triple bond have?
6. For the compound C2Cl2H2:
a. Make a Lewis dot structure of the non-polar isomer.
b. Draw the other two isomers.
c. Name the shape.
d. Which of the isomers have possibility for cis-trans orientation?
e. What is the orbital hybridization?
f. Draw an orbital box diagram for the non- polar isomer.
7. What is the strongest type of bonding? Explain your answer.
8. What is the weakest type of bonding? Explain your answer.
9. What are the characteristics of covalent bonds?
10. When do hydrogen bonds form?
11. What does the term electrolyte mean? Name the electrolytes as liquids.
12. List the substances in order of increasing polarity:
SO2 H2O C2H5OH CH4
13. Draw the shape of a 1. tetrahedral, 2. bent, 3. flat triangle molecule.
14. Match the molecule to the shape you drew above:
CH4___________
H2O___________
BCl3___________
15.Name the angle measures between the bonds from problem 13?
1. ________ 2. ________ 3. ________
16. When do substances dissolve; in other words, what kinds of substances dissolve each other?
17. Name an element that cannot form a double bond and explain why this is so.
Hydrogen (and Helium and all halogens) cannot form a double bond because it only has one free unpaired
electron.
18. Draw a Lewis dot structure for PCl5
a. Sketch the shape.
b. Name the shape.
c. Are the bonds polar?
d. Is the molecule polar?
19. What is the VSEPR theory?
20. Draw orbital box diagrams and Lewis dot structures for the following:
a. OF2
HCN
b. PH3
c.
Colligative Properties Review Problems
(Boiling and Freezing Point of Solutions)
Kf of water is 1.86 C/m.
Kb of water is .512oC/m
1. An aqueous solution of an unknown non-ionizing substance has a freezing point of -1.56°C.
What is the molality of the solution?
2. The molality of an aqueous solution of NaCl is .0250 M. What is the freezing point?
3. In the solution from problem 2, what is the change in boiling point?
4. You have an aqueous solution of NaCl with a molality twice that of a solution of Ca(NO3)2. Which has a
lower freezing point?
5. A 2 molal solution of NaCl is diluted to 0.75 molal. What is the change in boiling point?
6. What is the freezing point of a solution containing 0.5 mol of glucose dissolved in 200g of water? (Kf for
water is 1.86oC/m)
7. To what volume must 10.0mL of 5.00 M HCl be diluted to make 0.500 M HCl solution?
8. Which of the following is the solution with a highest boiling point (lowest freezing point)?
A. 0.1 M MgCl2
B. 0.1 M HClO4
C. 0.1 M NH4OH
D. 0.1 M LiNO3
9. If you raise the temperature the solubility of solid solutes in water will increase or decrease?
10. Which of the following will be the most electrically conductive?
A. Sugar dissolved in water
B. Salt water
C. Salt dissolved in an organic solvent
D. An oil and water mixture
11. Calculate the molal boiling point elevation of 11.00 g. aqueous solution with 0.0230 mol. of methanol
(CH3OH). (Kb for water=0.512°C/m)
12. Calculate the mass of NaCl needed to depress the freezing point of a 400.0 g. aqueous solution by 2.34°C
(Kf for water=1.84°C/m; d of NaCl=2.16 g/cm3)
13. By how much is the vapor pressure of a 100.00 g. aqueous solution containing 35.00 g. of KI at 25°C
lowered?
14. What mass of solute is needed in order to lower the vapor pressure of a 45.00 g. aqueous sol-n of KI at
30°C by 5.00 mmHg
15. Calculate the molality of an aqueous solution of NaCl if it has a freezing point depression of 1.00°C and a
kb of 1.84°C/m
Answers (Colligative Properties)
1. .839 mol/kg
2. -.930 C
3. +.0256 C
4. NaCl solution
5. -1.30 C
6. To solve this problem you need to find the molality of the solution:
Molality = moles of solute/ kg of solvent = 0.5 mol/ 0.200 kg = 2.5m
Using the equation ΔTf =Kf mn you find the freezing point depression:
2.5m x 1.86 oC/m = 4.65 oC ( n=1, glucose does not ionize)
0- 4.65 = - 4.65 oC
7. Use the equation McVc = MdVd (Mc = molarity of a concentrated solution, Md = molarity of a diluted
solution, Vc = volume of a concentrated solution, Vd = volume of a diluted solution
10.0 mL = 0.0100 L
(5.0 M)(0.0100L) = (0.50M)( Vd)
Vd = 0.100L = 100 mL
8. To answer this question you have to know that the greater the number of dissolved particles, the
greater is the boiling point elevation or the freezing point depression is. Since the molar concentration
for all the solutions is the same, therefore you just have to choose the solution that has the most
particles.
So MgCl2 breaks into 3 particles, while all of the other solutions break only into 2 particles. MgCl2 has
the highest boiling point and the lowest freezing point. The answer is A.
9. Increase. A solid solute is more soluble at higher temperatures and less soluble at lower temperatures
10. …B… Salt is an ionic substance, so when dissolved it will become an electrolyte and, therefore,
conduct electricity.
32.06
g.
CH
3
OH
11. 0.0230 mol CH3OH 
= 0.737g. CH3OH
1
mol
CH
3
OH
Mass of solvent=Mass of sol-n – Mass of solute=10.26g. sol-n
1kg.
10.6 g. sol-n  1000g. = 0.0106kg.
∆T=Kb*m
∆T=2.17m*0.512/m=1.11ºC
0
.0230
mol
CH
3
OH
m= 0.0106
kg.
sol
-n= 2.17m
12. ∆T=2.34 ºC
m * Kf =2.34 ºC
1.84ºC/m * m=2.34 ºC
m=1.27m
mol
solute
58.44g.
m


mol
solute
= mass of solute
kg
solvent
1mol.
mol solute=mass of solute/58.44 g.
let mass of solute=x
then g. of solvent= 400.0-x
1kg. 400
.0 x
kg.
and kg. of solvent=(400.0-x)g.  1000g. =
1000
x g.
17.11x
58.44g./mol.
=1.27m
=1.27m
400.0- x
400 - x
kg.
1000
17.11x=508-1.27x
18.38x=508
x=27.64 g.
13. ∆P=xsolute*Pº
Pº=23.76 mmHg
1mol.
KI
mol solute=35.00g. KI 166
.8g.KI=0.2098 mol KI
1
mol
H
2O
mol solvent=65.00 g. H2O 18.02
=3.610 mol. H2O
g.
H
2O
mol sol-n=mol solute + mol solvent=3.820 mol sol-n
xsolute=0.2098 mol. KI/3.820 mol sol-n=0.05492
∆P=23.76mmHg*0.05492=1.305mmHg
14. ∆P=Pº*xsolute
Pº=31.8mmHg
∆P=5.00mmHg
5.00mmHg=31.8mmHg* xsolute
xsolute=0.157
mol solute*166.0g/mol=mass of solute
mass
ofsolute
mol solute= 166.0
g/mol
45
.
00
g.
-mass
of
solute
mol solvent=
18.02
g/mol
mass
ofsolute
45
.
00
g.
-mass
of
solute
7470
.
g.

147
.
98
m
ass
of
solute
mol sol-n= 166
+
=
.0g/mol
18.02
g/mol 166.0

18.02
g/mo
l
m
ass
of
solute
166.0g/mol
0.157= 7470.g.
-147.98mass
of
solute
18.02

166.0
g/mo
l
mass
of
solute
147.98mass
of
solute
0.157= 7470.g.
18.02g/mol
7470
.
g
.
147.98mass
of
solute
18.02
115 mass of solute=7470.g.-147.98mass of solute
mass of solute=0.157
263mass of solute=7470g.
Mass of solute=28.4 g.
15. ∆T=Kb*m
1.00ºC=1.86 ºC/m*m
m=0.538m