General Chemistry II: Reactions in Aqueous Solution, Stoichiometry, and Solution Properties – Study Guide

Reactions in Aqueous Solution

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water. These rules are essential for determining the products of precipitation reactions and for writing net ionic equations.

• Most nitrate (NO3-) salts are soluble.

• Most salts containing alkali metal ions (Li+, Na+, K+, etc.) and ammonium (NH4+) are soluble.

• Most chloride, bromide, and iodide salts are soluble, except those of Ag+, Pb2+, and Hg22+.

• Most sulfate (SO42-) salts are soluble, except BaSO4, PbSO4, Hg2SO4, and CaSO4.

• Most hydroxide (OH-) salts are only slightly soluble, except those of alkali metals and Ba(OH)2, Sr(OH)2, Ca(OH)2 (slightly soluble).

• Most sulfide (S2-), carbonate (CO32-), chromate (CrO42-), and phosphate (PO43-) salts are only slightly soluble, except those containing alkali metals or NH4+.

Example: Mixing solutions of AgNO3 and NaCl produces a precipitate of AgCl, since AgCl is insoluble.

Activity Series of Metals

The activity series ranks metals by their ability to be oxidized (lose electrons). A metal higher in the series will displace a metal ion lower in the series from solution.

Metal	Oxidation Reaction
Lithium	Li(s) → Li+(aq) + e-
Potassium	K(s) → K+(aq) + e-
Calcium	Ca(s) → Ca2+(aq) + 2e-
Sodium	Na(s) → Na+(aq) + e-
Magnesium	Mg(s) → Mg2+(aq) + 2e-
Aluminum	Al(s) → Al3+(aq) + 3e-
Zinc	Zn(s) → Zn2+(aq) + 2e-
Iron	Fe(s) → Fe2+(aq) + 2e-
Lead	Pb(s) → Pb2+(aq) + 2e-
Hydrogen	2H+(aq) + 2e- → H2(g)
Copper	Cu(s) → Cu2+(aq) + 2e-
Silver	Ag(s) → Ag+(aq) + e-
Gold	Au(s) → Au3+(aq) + 3e-

Example: Zinc metal will displace copper(II) ions from solution, but copper will not displace zinc ions.

Stoichiometry and Solution Calculations

Concentration and Dilution

Molarity (M) is the number of moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution: $M_1V_1 = M_2V_2$ (where 1 = initial, 2 = final)

Example: To prepare 0.1 M NaCl from 1.0 M NaCl, use $V_1 = \frac{M_2V_2}{M_1}$.

Limiting Reactant and Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: The maximum amount of product possible, calculated from the limiting reactant.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Example: If 6.000 g of Cu2S is recovered from a reaction with a theoretical yield of 7.000 g, percent yield is $\frac{6.000}{7.000} \times 100\% = 85.7\%$.

Balancing Redox Reactions

Half-Reaction Method

Redox reactions can be balanced by separating them into oxidation and reduction half-reactions, balancing each for mass and charge, and then combining them.

1. Write the oxidation and reduction half-reactions.

2. Balance all elements except H and O.

3. Balance O by adding H2O; balance H by adding H+.

4. Balance charge by adding electrons.

5. Multiply half-reactions to equalize electrons, then add together.

Example: For Zn(s) + Ag+(aq) → Zn2+(aq) + Ag(s):

• Oxidation: Zn(s) → Zn2+(aq) + 2e-

• Reduction: 2Ag+(aq) + 2e- → 2Ag(s)

• Overall: Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)

Gas Laws and Stoichiometry

Ideal Gas Law

The ideal gas law relates pressure, volume, temperature, and moles of a gas.

• Equation: $PV = nRT$

• P = pressure (atm), V = volume (L), n = moles, R = 0.08206 L·atm/(mol·K), T = temperature (K)

Example: To find the number of moles of H2 gas collected over water, correct for vapor pressure of water and use the ideal gas law.

Useful Integrals and Constants

Integral	Result
\( \int x^n dx \)	\( \frac{x^{n+1}}{n+1} + C \)
\( \int \frac{dx}{x} \)	\( \ln|x| + C \)
\( \int e^{ax} dx \)	\( \frac{1}{a}e^{ax} + C \)

Gas Constant: R = 0.08206 L·atm/(mol·K)

Standard Temperature and Pressure (STP): 0°C (273.15 K), 1 atm

Periodic Table and Properties

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties. It is essential for predicting element behavior, valence, and reactivity.

• Groups: Vertical columns; elements in the same group have similar valence electron configurations.

• Periods: Horizontal rows; properties change progressively across a period.

Example: Alkali metals (Group 1) are highly reactive and form +1 ions.

Types of Chemical Reactions in Aqueous Solution

• Precipitation Reactions: Formation of an insoluble product (precipitate) from two soluble reactants.

• Acid-Base Reactions: Transfer of protons (H+) between reactants.

• Redox Reactions: Transfer of electrons between species; involves changes in oxidation states.

• Gas-Forming Reactions: Reactions that produce a gas as a product (e.g., CO2, H2).

Example: Mixing Na2CO3 and HCl produces CO2 gas.

Net Ionic Equations

Net ionic equations show only the species that actually change during the reaction, omitting spectator ions.

• Write the balanced molecular equation.

• Write the complete ionic equation (all strong electrolytes as ions).

• Cancel spectator ions to get the net ionic equation.

Example: For AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq):

• Net ionic: Ag+(aq) + Cl-(aq) → AgCl(s)

Summary Table: Key Equations and Constants

Equation/Constant	Value/Expression
Ideal Gas Law	$PV = nRT$
Molarity	$M = \frac{\text{mol solute}}{\text{L solution}}$
Dilution	$M_1V_1 = M_2V_2$
Percent Yield	$\frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$
Gas Constant (R)	0.08206 L·atm/(mol·K)
STP	0°C (273.15 K), 1 atm