Chapter 4 - Types of Chemical Reactions and Solution Stoichiometry
AP Chemistry
Water, the Common Solvent
• Water behaves as a great solvent for many substances due to its unique molecular structure.
– Solvent – dissolving medium in a solution
– Solute – a substance dissolved in a liquid to form a solution
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Molecular Structure of Water
– H2O
– bent molecular geometry
– ~105o between each O-H bond
– polar molecule
Polarity and Hydration
• It is polarity that gives water its great ability to dissolve substances.
• Hydration – interaction between solute particles and water molecules
– positive ends of water molecule attract negative ions or negative ends of other polar molecule
– negative ends of water molecule attract positive ions or positive ends of other polar molecules
Hydration and Solubility
• δ+ H attracts Cl- ions
• δ- O attracts Na+ ions
• A similar scheme can be used to illustrate the dissolution of a molecular substance in water.
• “Like dissolves like”
The Nature of Aqueous Solutions: Strong and Weak Electrolytes
• Solutions, homogeneous mixtures consisting of a solute and a solvent, can be characterized by their ability to
conduct an electric current.
• Electrical Conduction – phenomenon in which a material contains movable particles with electric charge
• Different solutions can have different electric conductivities.
Strong Electrolytes
• Strong electrolytes
– Carry electric current well
– completely ionize in water
– Include soluble salts, strong acids and strong bases
• Example:
HCl(aq)  H+(aq) + Cl-(aq) ~100%
Weak Electrolyte
• Weak electrolytes
– Carry electric current poorly
– Small degree of ionization in water
– Include slightly soluble salts, weak acids and weak bases
• Example:
NH3(aq)  NH4+(aq) + OH-(aq) ~1%
Nonelectrolytes
• Nonelectrolytes
– Carry no electric current
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Do not produce ions in water
Include many soluble organic compounds like alcohols and sugars
The Composition of Solutions
• Solutions consist of solutes dissolved in solvents
• Solution concentration can vary depends on the amount of each substance present
Molarity
• Molarity – moles of solute per volume of solution in liters
•
M = molarity = moles of solute
liter of solution
Molarity Example
• What is the molarity of a solution prepared by dissolving 25.0 g of solid NaCl in enough water to make a 2.0 L
solution?
Dilutions
• Dilution – addition of solvent to a solution in order to achieve a desired molarity
• Dilution with water does not alter the number of moles of solute present.
• M1V1 = M2V2
• M1 and V1 correspond to initial molarity and volume, and M2 and V2 correspond to the final molarity and volume
Dilution Example
• How many milliliters of a 12.0 M concentrated solution of hydrochloric acid are needed to produce 5.00 L of a
1.00 M HCl solution?
Types of Reactions
• Virtually all reactions can be classified into the following three groups:
– Precipitation reactions
– Acid-base reactions
– Oxidation-reduction reactions
Precipitation Reactions
• Precipitation Reactions – two solutions are mixed to form an insoluble product that separates from solution
• Precipitate – solid that forms in a precipitation reaction
Predicting Precipitates
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Precipitates can be predicted by examining the ions present in solution.
If any ion pair collisions produce insoluble products, then a precipitate will form.
Example:
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
Solubility Rules
1. Most nitrate and nitrite salts are soluble.
2. Most salts containing the alkali metals and ammonium ion are soluble.
3. Most chloride, bromide, and iodide salts are soluble. (Exceptions: Ag+, Pb2+, and Hg22+)
4. Most sulfate salts are soluble. (Exceptions: Ba2+, Pb2+, Ca2+, Hg22+)
5. Most hydroxide salts are slightly soluble. Soluble hydroxides include NaOH and KOH. Ba(OH)2, Sr(OH)2, and
Ca(OH)2 are slightly soluble.
6. Most sulfide, carbonate, chromate, and phosphate salts are only slightly soluble.
Describing Reactions in Solution
• Reactions in solutions can be described using the following equations:
1. Formula– gives overall reaction stoichiometry
2. Ionic – represents as ions all reactants and products that are strong electrolytes
3. Net-ionic – includes only those components undergoing a change, spectator ions are not included
Example
• Write the molecular, ionic, and net-ionic equation describing the reaction between solutions of lead (II) nitrate
and potassium chloride.
• Molecular Equation:
Pb(NO3)2(aq) + 2KCl(aq)  PbCl2(s) + 2KNO3(aq)
• Ionic Equation:
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2Cl-(aq)  PbCl2(s) + 2K+(aq) + 2NO3-(aq)
• Net-Ionic Equation:
Pb2+(aq) + 2Cl-(aq)  PbCl2(s)
Stoichiometry of Precipitation Reactions
1. Identify the species present in the combined solution, and determine what reaction occurs.
2. Write the balanced chemical equation.
3. Calculate moles of the reactants.
4. Determine which reactant is limiting.
5. Calculate the moles of product, or products.
6. Convert to required units.
Stoichiometry Example I
• When potassium iodide and lead (II) nitrate are mixed together a bright yellow PbI2 precipitate forms. Calculate
the mass of PbI2 that forms when 10.0 mL of 0.100 M KI is mixed with 20.0 mL of 0.0750 M Pb(NO3)2.
Stoichiometry Example II
• What volume of 0.100 M Na2SO4 is required to precipitate all the lead(II) ions from 150.0 mL of 0.250 M
Pb(NO3)2?
Acid-Base Reactions
• Acid-Base Reactions are also known as acid-base neutralization reactions.
• Brönsted-Lowry Definition:
– An acid is a proton donor
– A base is a proton acceptor
• There are other definitions of acids and bases which will be discussed in detail in chapter 14.
Predicting Acid-Base Reactions
• Acid + Base → Salt + Water
• Strong Acid/Strong Base
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
Net-Ionic: H+(aq) + OH-(aq)  H2O(l)
• Weak Acid/Strong Base
CH3COOH(aq) + KOH(aq)  H2O(l) + CH3COOK(aq)
Net-Ionic: CH3COOH(aq) + OH-(aq)  H2O(l) + CH3COO-(aq)
Acid-Base Stoichiometry
• Step 1: Determine whether or not the species present will react.
• Step 2: Write the balanced net ionic equation for the reaction.
• Step 3: Calculate the moles of reactants. Use molarities if reactants are in solution.
• Step 4: Determine the limiting reagent where appropriate.
• Step 5: Calculate the moles of the required reactant or product.
• Step 6: Convert to the appropriate units.
Example
• What volume of 0.500 M HCl solution is needed to neutralize 50.0 mL of 0.200 M NaOH?
Acid-Base Titrations
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Titrant
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Volumetric Analysis – technique used for determining the amount of a certain substance by doing a titration
Titration – delivery of a measured volume of a solution of known concentration (the titrant) into a solution
containing the substance being analyzed (the analyte)
Titrant – solution of standardized concentration
A titrant must react with analyte in a known manner. (Acids and bases undergo neutralization reactions)
Enough titrant must be added to react exactly with the analyte, this is called the endpoint or equivalence point.
Indicator – substance added at the beginning of the titration that changes color at the equivalence point
For a Successful Titration:
1. The reaction between titrant and analyte must be known.
2. The equivalence point must be marked accurately.
3. The volume of titrant added to reach equivalence point must be known.
Acid-Base Titration Example
• 37.0 mL of standardized 0.100 M NaOH is used to titrate a hydrochloric acid solution of unknown concentration.
How many moles of HCl are present in the unknown acid solution? What was the molarity of the original HCl
solution if there were 10.0 mL present before titration?
Oxidation-Reduction Reactions
• Reactions in which one or more electrons are transferred are called oxidation-reduction reactions.
• Examples:
Mg(s) + S(s)  MgS(s)
CH4(g) + O2(g)  CO2(g) + 2H2O(l) + E
Oxidation States
• Imaginary charges atoms would have if covalently bound atoms transferred electrons or actual charges on ions
in ionic compounds. (Assume the more electronegative atom takes on negative charge)
• Examples:
CO2: C=4+, O=2O2: O = 0 (equally shared)
PCl5: P=5+, Cl=1Rules for Assigning Oxidation States
1. An atom in elemental form receives an oxidation state of 0. (O2, Na, Fe, Cl2 = 0)
2. Oxidation states for monatomic ions are the same as their charge. (Na+=1+, O2-=2-)
3. In its compounds, fluorine always has an oxidation state of 1-. (HF, F=1-)
4. In covalent compounds with nonmetals, hydrogen is assigned an oxidation state of 1+. (HCl, NH3, H=1+)
5. Oxygen is usually assigned 2-. Exceptions include peroxides (O2 group O2=2-) and compounds containing fluorine
(OF2. O=2+).
6. Sum of the oxidation numbers must equal zero, unless substance is an ion.
Characteristics of Redox Reactions
• Oxidation and reduction occur simultaneously in a redox reaction.
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– If a species is being oxidized, another species must be reduced.
Oxidation – increase in oxidation state
(loss of electrons)
Reduction – decrease in oxidation state
(gain of electrons)
Oxidizing and Reducing Agents
• Oxidizing Agent – substance being reduced
• Reducing Agent – substance being oxidized
Summary
• Oxidized
– Loses electrons
– Oxidation state increases
– Reducing agent
– OIL: Oxidation is losing
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Reduced
– Gains electrons
– Oxidation state decreases
– Oxidizing agent
– RIG: Reduction is gaining
Example
• Identify the oxidizing and reducing agent in the following redox reaction. 2Al(s) + 3I2(s)  2AlI3(s)
• Al: Al0  Al3+, loses e-, is oxidized, and is the reducing agent
• I2: I20  I-, gains e-, is reduced, and is the oxidizing agent