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Chem 1A
Dr. White
Handout 4
4.4 Types of Chemical Reactions (Overview)
A. Non-Redox Rxns
B. Oxidation-Reduction (Redox) reactions
4.6. Describing Chemical Reactions in Solution
A. Molecular Equation
B. Complete Ionic Equation
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C. Net Ionic Equation
4.9. Oxidation-Reduction Reactions (“Redox Reactions”)
A. Oxidation:
B. Reduction:
C. Oxidizing Agent:
D. Reducing Agent:
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3 E. Types of Redox Reactions
1. Combination (Composition)
Metal + Oxygen  Metal Oxide OR Nonmetal + Oxygen  Nonmetal Oxide
Example: Carbon + limited Oxygen
and Carbon + excess Oxygen
2. Displacement
a. Metal A + Metal B Salt  Metal A Salt + Metal B
Example: Magnesium + aqueous silver nitrate
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4 b. Active Halogen + Metal Halogen Salt  Less Active Halogen + Metal
Halogen Salt
Example: Fluorine + Aqueous Sodium Chloride
c. Metal + Acid  Metal salt + hydrogen gas
Example: Aluminum + Hydrochloric acid
d. Metal + Water  Metal hydroxide + hydrogen gas
Example: Lithium + Water
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3. Combustion – reactions with oxygen
one common type: Hydrocarbon + Oxygen → CO2 (g) + H2O (l or g)
Example: C4H10 (g) + oxygen
F. Oxidation Numbers
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6 Lecture Examples:
a. What is the oxidation state of chromium in Cr2+?
b. What is the oxidation state of chromium in CrCl3?
c. What is the oxidation state of chromium in the dichromate ion?
M M+ e-­‐ X X-­‐ Oxidized: Loses electrons Oxidation # increases Reducing agent Reduced: Gains electrons Oxidation # decreases Oxidizing agent Lecture Examples: In each of the following examples, determine has been
oxidized and what has been reduced and which is the oxidizing agent and which
is the reducing agent.
a. Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)
b. PbO (s) + CO (g) → Pb (s) + CO2 (g)
c. sodium metal reacts with water.
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18.1 Balancing Redox Reactions
A. Balancing by Inspection
1. Lecture Example: Balance the following redox reaction:
Zn (s) + Ag+ (aq) → Zn2+ (aq) + Ag (s)
B. Balancing by the Half Reaction Method
1. Balancing redox reactions that occurs in an acidic solution:
Example: Cr2O72- + NO2- → Cr3+ + NO3Step 1: Split into half reactions
Step 2: Balance atoms and charges in the half reactions (MOHE)
M: balance miscellaneous atoms (atoms other than O and H)
O: balance O by adding water molecules
H: balance H by adding H+ ions
E: balance charges by adding electrons
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Step 3: Multiply reactions by an integer so that the number of electrons are
equal
Step 4: Add half reactions together
Step 5: Check that the atoms and charges are balanced
2. Balancing redox reactions that occurs in a basic solution:
Example: MnO4- (aq) + SO32- (aq) → MnO2 (s) + SO42- (aq) (basic)
Follow steps 1-5:
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9 Step 6: Note the number of H+ ions in the equation, add this number of OH- ions to
both sides
Step 7: Simplify by noting that H+ reacts with OH- to give H2O
3. Balance the following: Cr(OH)3(s) + ClO3-( aq) --> CrO42-(aq) + Cl-(aq)
(basic)
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10 4.5 Precipitation Reactions
A. Precipitation Reactions Overview
Solubility rules
1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this
rule are rare. Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally soluble.
3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are
halide salts of Ag+, Pb2+, and Hg22+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.
4. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4,
Ag2SO4, Hg2SO4 and CaSO4.
5. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are
soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide
salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not
soluble.
6. Most sulfides, carbonates, chromates, and phosphates are only slightly soluble except for
those containing alkali metals and NH4+.
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11 Lecture Examples:
a. Determine the net ionic equation for the following reaction:
Ba(NO3)2 + Na2SO4 → BaSO4 + 2NaNO3
b. Write molecular and net ionic equations for the following:
i. aqueous solutions of sodium chloride and iron (II) nitrate are
mixed
ii. Aqueous solutions of aluminum sulfate and sodium hydroxide
are mixed
4.8 Acid-Base Reactions
A. Acid-Base Definitions
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B. Acid-Base reactions
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13 3.9 Stoichiometric Calculations
A. Mole to Mole Conversions
Lecture Example: How many moles of hydrogen gas will form if 55.61 mol
of aluminum completely reacts with an excess of HCl?
B. Gram to gram conversions
Lecture Example: 96.1 g C3H6 reacts with oxygen to form carbon dioxide
and water. How much oxygen do we need to react with all the C3H6?
Lecture Example: Baking soda (NaHCO3) is used as an antacid. It
neutralizes HCl in the stomach. How many grams of HCl are neutralized
per 1.00 gram of baking soda? If all the baking soda reacts, how many
grams of water is produced?
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14 3.10 Calculations involving Limiting Reactants
A. Limiting reactants
A simple example:
Lecture Example: Carbon disulfide (CS2) burns in oxygen according to the
following equation. CS2 + 3 O2 → CO2 + 2 SO2
Calculate the moles of SO2 each component present in the flask at the end
of the reaction when 3.0 mol of CS2 and 3.0 mol of O2 are mixed.
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15 Lecture Example: Nitrogen gas is prepared by passing ammonia gas
over solid copper (II) oxide at high temperatures. The other products are
solid copper and water vapor. If 18.1 g of NH3 are reacted with 90.4 g of
CuO, which is the limiting reactant? How many grams of nitrogen gas will
be formed? How much of the excess reactant is left?
Lecture Example: For the reaction below, if you start with 150.0 g Fe2O3 and
140.0 g CO, how many grams of Fe will you produce?
Fe2O3 (s) + 3 CO (g) → 2 Fe (s) + 3 CO2 (g)
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B. Theoretical Yield
C. Actual Yield
D. Percentage Yield
Lecture Example: if actual yield of Fe (s) produced from the previous
example is 98.9 g, what is the percent yield?
4.7 Stoichiometry of Precipitation Reactions
A. Lecture Examples:
1. How many grams of precipitate form when 35.0 mL of 0.160M barium
chloride reacts with 58.0 mL of 0.065 M sodium sulfate?
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17 2. 350.0 mL of 0.210M sodium sulfate reacts with 0.500 L of 0.196M barium
nitrate. A precipitate forms. What are the concentrations of all the ions
after the reaction is complete?
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18 3. 0.250 L of 0.100M silver nitrate is mixed with 0.500 L of 0.100M sodium
sulfate. A precipitate forms. What are the concentrations of all the ions
after the reaction is complete?
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19 4.8 Titrations
A. Types of Titration
B. Equivalence Point
C. Examples:
1. You perform an acid-base titration to standardize an HCl solution by
placing 50.00 mL in a flask with a few drops indicator solution. You put
0.1524 M NaOH in the buret with an initial buret reading of 0.55 mL. At
the end point the reading on the buret is 33.87 mL. What is the
molarity of the acid solution?
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20 2. What volume of 0.1292 M Ba(OH)2 would neutralize 50.00 mL of the HCl
solution standardized in the preceding example.
3. A common analysis for iron is titration of iron (II) with potassium
dichromate, K2Cr2O7. The half reactions associated with this reaction are
below:
Cr2O7-2 + 14H+1 +6e-  2Cr+3 + 7H2O
Fe+2  Fe+3 + eIf a 5.00 mL sample of an unknown solution that contains iron (II) and iron
(III) requires 28.45 mL of 0.100 M potassium dichromate to reach the
equivalence point, calculate the concentration of Fe2+ in the unknown
solution.
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21 4. Calcium ion is required for blood to clot and for many other cell
processes. An abnormal calcium ion concentration is indicative of disease.
To measure the calcium ion concentration, 1.00 mL of human blood is treated
with a Na2C2O4. The resulting precipitate is dissolved in dilute acid to
release C2O42- into solution and allow it to be oxidized using KMnO4. This
solution required 2.05 mL 4.88 x 10-4 M KMnO4 to reach the endpoint of the
titration. Calculate the molar concentration of calcium ion in the blood.
NOTE: One product of the rxn is Mn2+ and the other product of the redox
reaction is a gas that turned limewater cloudy.
3.5 Percent Composition of Compounds –
A. Mass Percent from the Chemical Formula
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22 Lecture Example: Consider magnesium chlorate.
a. What is the mass % of each atom in magnesium chlorate?
b. How many grams of oxygen are in 16.55 g of magnesium chlorate?
Lecture Example: A farmer determines that 60. Lbs N per acre is
necessary in her field. If she is using ammonium nitrate as fertilizer, how
many pounds of ammonium nitrate must she spread on 1.0 acre?
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23 3.6 Determining the Formulas of Compounds
A. Determining formulas from data:
1. Determining formulas from mass % (we did this already)
Try this example on your own to review: During physical activity, lactic
acid (molar mass = 90.08 g/mol) forms in muscles and is responsible for
muscle soreness. Elemental analysis shows that this compound contains 40.0
mass % C and 53.3 mass % O. Determine the empirical formula and the
molecular formula for lactic acid.
2. Combustion analysis of organic compounds
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24 Lecture Example: A 0.589 g sample of an organic compound containing only
carbon, hydrogen and oxygen was burned completely in air to produce
0.733 g of CO2 and 0.299 g of H2O. What is the empirical formula of the
compound?
Chapter 5: Gases
5.2 The Gas Laws of Boyle, Charles, Avogadro
A. Boyle's Law
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B. Charles' Law
C. Avogadro’s Law
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26 5.3 The Ideal Gas Law
A. Derivation
B. Examples using the ideal gas law:
1. Lecture Example: 2.96g Mercuric Chloride vaporized in a 1.0L bulb at 680
K and P=453 torr. What is the molar mass?
2. What is the density (in g/L) of CO2 at 30.3°C and 744 mmHg?
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3.
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27 A balloon contains 5.41 dm3 of He at 24°C and 101.5 kPa. The gas in the
balloon is heated to 35°C and the pressure becomes 102.8 kPa. What is
the volume of the gas? (101.3 kPa = 1 atm)
5.4 Gas Stoichiometry
A. Lecture Example:
A sample of CH4(g) having a volume of 2.80 Liters at 25°C and 1.65 atm, was
ignited with a sample of oxygen gas having a volume of 35.0 Liters at 31°C
and 1.25 atm to produce CO2 and H2O vapor. Calculate the volume of CO2
formed at a pressure of 2.50 atm and a temperature of 125°C.
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28 Copper dispersed in absorbent beds is used to react with oxygen impurities
in ethylene used for producing polyethylene. The beds are regenerated
when hot H2 reduces the metal oxide forming the pure metal and H2O. On a
laboratory scale, what volume of H2 at 765 torr and 25°C is needed to
reduce 35.5 g of copper (II) oxide?
The Group IA metals react with the halogens (Group 7A) to form ionic metal
halides. What mass of potassium chloride forms when 5.25 L of chlorine gas
at 0.950 atm and 293 K reacts with 17.0 g of potassium?