4.4 Types of Chemical Reactions (Overview) A. Non …€¦ · · 2014-12-04Describing Chemical...

Chem 1A Dr. White Fall 2014 1

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


C. Net Ionic Equation

4.9. Oxidation-Reduction Reactions (“Redox Reactions”)

A. Oxidation:

B. Reduction:

C. Oxidizing Agent:

D. Reducing Agent:


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


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


3. Combustion – reactions with oxygen

one common type: Hydrocarbon + Oxygen → CO2 (g) + H2O (l or g)

Example: C4H10 (g) + oxygen

F. Oxidation Numbers


Oxidized: Loses electrons

Oxidation # increases

Reducing agent

Reduced: Gains electrons

Oxidation # decreases

Oxidizing agent

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?

e-‐

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.

M M+

X X-‐


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+ + NO3-

Step 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


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) + SO4

2- (aq) (basic) Follow steps 1-5:


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)


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 Hg2

2+. 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

+.


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


B. Acid-Base reactions


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?


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.


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)


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?


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? What is the mass of the ppt? Assume the volumes are additive.


3. 0.250 L of 0.100M silver nitrate is mixed with 0.500 L of 0.100M sodium carbonate. A precipitate forms. What are the concentrations of all the ions after the reaction is complete? What is the mass of the ppt? Assume the volumes are additive.


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?


2. What volume of 0.1524 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 + e-

If 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.


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 C2O4

2- 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 (not from data!)


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?


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


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


B. Charles' Law

C. Avogadro’s Law


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?


3. 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.


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?

4.4 Types of Chemical Reactions (Overview) A. Non …€¦ · · 2014-12-04Describing Chemical...

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