Download AP Chemistry Chapter 4 Notes: Types of Chemical Reactions and

AP Chemistry
Chapter 4 Notes: Types of Chemical Reactions and Solution Stoichiometry
4.1Water, the Common Solvent
A. Structure of water
1. Oxygen's electronegativity is high (3.5) and hydrogen's is lower (2.1)
2. Water is a bent molecule with polar bonds
3. Water is a polar molecule
B. Hydration of Ionic Solute Molecules
1. Positive ions attracted to the oxygen end of water
2. Negative ions attracted to the hydrogen end of water
C. Hydration of Polar Solute Molecules
1. Negative end of polar solute molecules are attracted to water's hydrogen
(with its partial positive charge)
2. Positive end of polar solute molecules are attracted to water's oxygen
D. "Like Dissolves Like"
1. Polar and ionic compounds dissolve in polar solvents like water
2. Nonpolar compounds like fats dissolve in nonpolar solvents.
4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes
A. Definition of Electrolyte: a substance that when dissolved in water produces a
solution that can conduct an electric current
B. Strong electrolytes conduct current very efficiently since they completely ionize
when dissolved in water
1. Ionic compounds
2. Strong acids (HNO3(aq), H2SO4(aq), HCl(aq))
3. Strong bases (KOH, NaOH)
C. Weak electrolytes conduct small currents since they only slightly ionize in water.
1. Weak acids (organic acids): acetic, citric, ...
2. Weak bases (ammonia...)
D. Nonelectrolytes conduct no current since no ions are
present in solution.
alcohols, sugars
4.3 The Composition of Solutions
A. Molarity (M)
Moles of solute / liters of solution
B. Concentration of ions in solution
Ionic compounds dissociate in solution, multiplying the molarity by the
number of ions present in one formula unit of the compound
C. Moles from concentration
Liters of solution x Molarity = moles of solute
D. Solutions of known concentration
1. Standard solution - a solution whose concentration is accurately known
2. Preparation of Standard solutions
L solution x mol/L x g/mol = grams required to prepare the standard
E. Dilution
1. Diluting a sample of solution with water does not change the number of
moles of solute in solution.
2. So, we use M1∙V1 = M2∙V2 to solve dilution problems since moles = moles
4.4 Types of Chemical Reactions
A. Precipitation reactions
When two solutions are mixed, an insoluble solid forms
B. Acid-Base reactions
A soluble hydroxide and a soluble acid react to form water and a salt
C. Oxidation-Reduction reactions (redox rxns)
Reactions in which one or more electrons are transferred
4.5 Precipitation Reactions
A. Dissociation
Ionic compounds dissolve in water and the ions separate from one another
AgNO3(aq) + NaCl(aq) products
Ag1+ (aq) + NO31−(aq) + Na+(aq) + Cl1−(aq)  products
B. Determination of Products
1. Recombination of ions
AgNO3 + NaCl AgCl + NaNO3
2. Elimination of reactants as products
AgNO3 and NaCl are reactants and can't be products
3. Identifying the precipitate
a. Reactants "switch partners" to form possible products.
b. AgCl and NaNO3 are the products
c. AgCl is insoluble, so it is the white precipitate.
d. If there is no insoluble product, the reaction does not occur
a. NaCl(aq) + KNO3(aq)  NaNO3(aq) + KCl(aq)
Both products are soluble and all ions remain independent in
solution; no reaction occurs: Na1+ (aq) + Cl1−(aq) + K+(aq) + NO31−(aq)
C. Solubility Rules for Salts in Water (p.148)
Soluble:
1. Salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the
ammonium (NH41+) ion are soluble.
2. Nitrate (NO31−) salts are soluble. Acetates are as well except with Ag+.
3. Halide (Cl−, Br−, and I−) salts are soluble. Exceptions are salts of Ag+,
Pb2+, and Hg22+.
4. Sulfate salts are soluble. Notable exceptions are Ag2SO4, BaSO4, PbSO4,
Hg2SO4 and CaSO4.
Insoluble:
5. Hydroxide salts are only slightly soluble. Exceptions: salts of Group I and
NH41+. Hydroxides of Ba2+, Sr2+, and Ca2+ are marginally soluble.
6. Sulfide (S2−), carbonate (CO32−), chromate (CrO42−), and phosphate
(PO43−) salts are only slightly soluble, except when combined with Group
I and NH41+.
4.6 Describing Reactions in Solution
A. Formula Equation
Written as a double displacement reaction rather than in the forms of
separated ions that actually exist in solution.
Na2CO3(aq) + Ca(NO3)2(aq)  2NaNO3(aq) + CaCO3(s)
B. Complete Ionic Equation
Represents all reactants and products as ions if they are strong electrolytes
2Na1+(aq) +CO32−(aq) + Ca2+(aq) +2NO31−(aq)  2Na1+(aq) + 2NO31−(aq) +CaCO3(s)
C. Net Ionic Equation
Includes only those components that take part in the chemical change.
Spectator ions are eliminated.
Ca2+(aq) + CO32−(aq)  CaCO3(s)
4.7 Stoichiometry of Precipitation Reactions
A. Determine what reaction takes place (formula and/or complete ionic equations)
B. Write the balanced net ionic equation.
C. Calculate the moles of reactants.
D. Determine which reactant is the limiting reagent.
E. Calculate the moles of product(s).
F. Convert to grams or other units, as required.
4.8 Acid-Base Reactions (Neutralization Reactions)
A. Definitions
1. Arrhenius: Acids produce H+ ions in solution, bases produce OH− ions.
2. Brønsted-Lowry: Acids are proton donors, bases are proton acceptors.
B. Net ionic equation for acid-base reactions
H+(aq) + OH−(aq)  H2O (l)
C. Stoichiometry calculations for acid-base reactions
1. List the species present in the combined solution as you do for a
precipitation reaction (including physical state symbols).
2. Write the balanced net ionic equation.
3. Calculate the moles of the reactants
For reactants in solution, use volumes (L) times their molarities (M).
4. Determine the limiting reactant where appropriate
5. Calculate the moles of the required reactant or product
6. Convert to grams or volume of solution as required
D. Acid-Base titrations
1. Titrant - Solution of known concentration (strong acid or strong base)
2. Analyte - Solution of unknown concentration
3. Equivalence point - Point at which the amount of titrant added to analyte
results in neutralization
4. Indicator - a substance added at the beginning of the titration that
changes color at the equivalence point
5. Endpoint - the point in the titration at which the indicator changes color
4.9 Oxidation-Reduction Reactions (Redox)
A. Electron transfer (OIL RIG)
1. Oxidation: Lose electrons (This substance is the reducing agent)
2. Reduction: Gain electrons (This substance is the oxidizing agent)
B. Rules for Assigning Oxidation Numbers
1. Oxidation number of an atom in an uncombined element = 0.
Includes polyatomic elements: diatomic molecules, allotropes (P4, S8...)
2. Oxidation numbers of monatomic (1 atom) ions = their charge. (Cl− = −1)
3. The sum of a compound's oxidation numbers = its charge.
Molecules = 0, Polyatomic ions = their charge
4. Assign oxidation numbers to the extremes in electronegativities first.
Follow this order keeping rule #3 in mind:
a) F = −1
b) Group 1A = +1 Exception: H = −1 if functioning as a hydride.
c) Group 2A = +2
d) Al = +3
e) O = −2
Exceptions: Peroxides (O22−) = −1; in OF2 = +2
f) Halogens = −1
Exception: if combined with more
electronegative halogens or oxygen.
C. Noninteger oxidation states
Some compounds have the same elements in different oxidation states:
1. Fe3O4 (magnetite)
a) Oxidation number for each iron averages to +8/3
b) Magnetite contains two Fe3+ ions and one Fe2+
c) Okay to mark oxidation state as +8/3
2. S4O62−
a) S = +5/2
b) two S's = 0; two S's = 5. Write as 5/2
4.10 Balancing Oxidation-Reduction Equations
A. Electron Exchange Method:
1. Assign oxidation numbers (to see which elements are changing)
+7 −2 −1 +1
+4 −2 0 +1−2
1−
−
+
MnO4 + I + H ➔ MnO2 + I2 + H2O
2. Connect elements that change oxidation numbers, indicating the number
of electrons gained or lost.
⌈‾‾‾‾‾‾‾‾‾‾‾+1‾‾‾‾‾‾‾‾‾‾⌉
MnO41− + I− + H+ ➔ MnO2 + I2 + H2O
⌊________−3 _____⌋
3. Check to see if there are any formulas involving more than one atom of
an element that changes oxidation states. Temporarily add a coefficient
that divides the formula by the subscript to make the compound
equivalent to one atom of that element.
... I− ... ➔ ... 1/2 I2 ...
[1/2 I2 = one atom of I]
4. Choose coefficients to balance the change in numbers of electrons, so
numbers of electrons lost = numbers gained on each side of the reaction.
⌈‾‾‾‾‾x3(+1)‾‾‾‾⌉
... 3I− ...➔ ... 3/2 I2 ...
This becomes the coefficient of the indicated elements.
5. Finish balancing the remaining equation.
MnO41− + 3 I− + 4H+ ➔ MnO2 + 3/2 I2 + 2H2O
6. Multiply the entire equation by any needed factor to remove fractions.
2 MnO41− + 6 I− + 8H+ ➔ 2 MnO2 + 3 I2 + 4H2O
Note: Some equations will not include water or water ions. Since these are
aqueous solutions, the water and/or water ions (hydroxide, hyronium)
may become part of the reaction. If this is the case:
7. Add necessary water or water ions to finish balancing the equation.
a) For acidic solutions:
1) add H2O to the side that needs O
2) add needed H+ to the other side to balance
b) For basic solutions:
1) add OH− to side that needs O (2 OH− for each O needed)
2) add H2O to finish balancing.
B. Half-Reaction Method
1. Assign oxidation numbers (as done in previous method step 1)
+7 −2 −1 +1
+4 −2 0 +1−2
MnO41− + I− + H+ ➔ MnO2 + I2 + H2O
2. Isolate the oxidation and reduction half-reactions. Balance as needed:
MnO41− ➔ MnO2
I− ➔ I2
3. Balance the numbers of atoms in the half reactions:
I− ➔ I2
becomes: 2 I− ➔ I2
4. Add water and/or water ions for needed O or H:
MnO41− ➔ MnO2 + 2 H2O
4H+ + MnO41− ➔ MnO2 + 2H2O
5. Add electrons to balance the change in oxidation number:
+7
−3 = +4
2(−1) = 0 − 2
+
1−
−
4H + MnO4 + 3e ➔ MnO2 + 2H2O
2I− ➔ I2 + 2e−
6. Choose coefficients to balance the gain / loss of electrons between the
half reactions:
2x [4H+ + MnO41− + 3e− ➔ MnO2 + 2H2O]
3x [2I− ➔ I2 + 2e−]
8H+ + 2MnO41− + 6e− ➔ 2MnO2 + 4H2O
6 I− ➔3 I2 + 6 e−
7. Add the two half reactions so all reactants are on one side, products on
the other:
8H+ + 2MnO41− + 6e− + 6 I− ➔ 2MnO2 + 4H2O + 3 I2 + 6 e−
8. Cancel equal quantities of electrons, H2O, H+, and OH−.
8H+ + 2MnO41− + 6 I− ➔ 2MnO2 + 4H2O + 3 I2