Chapter 5: Introduction to Solutions and Aqueous Reactions – Study Notes

Learning Objectives

• Define solution, solvent, solute, electrolyte, strong electrolyte, and weak electrolyte.

• Perform calculations involving molarity, solute mass or moles, and solution volume.

• Use dilution equations to solve for unknowns in dilution problems.

• Apply molar concentrations and stoichiometry to determine the amount of reactants or products in chemical reactions.

• Write total ionic and net ionic equations for molecular equations.

• Predict products and solubility in precipitation reactions using solubility rules.

• Define acid and base using the Arrhenius and Brønsted-Lowry definitions.

• Recognize and name strong acids and bases.

• Identify monoprotic, diprotic, and triprotic acids.

• Apply solution stoichiometry to titration and equivalence point calculations.

• Recognize gas-forming reactions and predict their products.

• Define and identify oxidation, reduction, oxidizing agent, and reducing agent.

• Assign oxidation numbers to elements in compounds and ions.

Solutions and Homogeneous Mixtures

Definition and Properties

A solution is a homogeneous mixture of two or more substances. In a solution, the composition is uniform throughout, and the individual components cannot be distinguished by the naked eye.

• The solvent is the component present in the greatest amount.

• The solute is the component present in a lesser amount and is dissolved by the solvent.

• A solution in which water is the solvent is called an aqueous solution.

• Example: When table salt (NaCl) is mixed with water, it forms a homogeneous solution. The salt appears to disappear, but it can be recovered by evaporating the water.

What Happens When a Solute Dissolves

Dissociation of Ionic Compounds

When an ionic compound dissolves in water, the water molecules surround and separate the individual ions from the crystal lattice. This process is called dissociation.

• Water molecules are polar and can stabilize ions in solution by surrounding them (solvation).

• For example, when NaCl dissolves, it separates into Na+ and Cl- ions.

• The resulting solution contains free-moving ions, which allows it to conduct electricity.

Table Salt Dissolving in Water

Microscopic View and Conductivity

• Each ion (Na+ or Cl-) is attracted to the surrounding water molecules and is pulled away from the crystal.

• Once in solution, the ions are surrounded by water molecules, which insulate them from other ions.

• This results in a solution with free-moving charged particles, making it an electrolyte capable of conducting electricity.

Electrolytes and Nonelectrolytes

Definitions and Examples

• An electrolyte is a substance that dissociates into ions when dissolved in water, producing a solution that conducts electricity.

• A strong electrolyte dissociates completely in water (e.g., NaCl, HCl).

• A weak electrolyte only partially dissociates in water (e.g., acetic acid).

• A nonelectrolyte dissolves in water but does not produce ions (e.g., sugar, methanol).

Solubility of Ionic Compounds

Soluble vs. Insoluble Compounds

• Soluble compounds dissolve well in water (e.g., NaCl).

• Insoluble compounds do not dissolve appreciably in water (e.g., AgCl).

• When an ionic compound dissolves, it breaks up into its constituent ions:

$\mathrm{NaCl_{(aq)} \rightarrow Na^+_{(aq)} + Cl^-_{(aq)}}$ $\mathrm{BaCl_2_{(aq)} \rightarrow Ba^{2+}_{(aq)} + 2Cl^-_{(aq)}}$ $\mathrm{Ca(NO_3)_2_{(aq)} \rightarrow Ca^{2+}_{(aq)} + 2NO_3^-_{(aq)}}$

• Solubility is often indicated by (aq) for aqueous (soluble) and (s) for solid (insoluble).

Solubility Rules for Ionic Compounds in Water

Solubility rules help predict whether an ionic compound will dissolve in water. The following table summarizes the main rules:

Ion/Compound	Solubility
Group 1A ions (Li+, Na+, K+, etc.)	Soluble
NH4+ (ammonium)	Soluble
NO3- (nitrate), C2H3O2- (acetate)	Soluble
Cl-, Br-, I-	Soluble (except with Ag+, Pb2+, Hg22+)
SO42- (sulfate)	Soluble (except with Ba2+, Pb2+, Ca2+, Sr2+)
CO32-, PO43-, OH-, S2-	Insoluble (except with Group 1A and NH4+)

Additional info: These rules are a summary; exceptions may exist for some transition metals and other ions.

Solution Concentration

Molarity

Molarity (M) is the most common unit of concentration in chemistry. It is defined as the number of moles of solute per liter of solution:

$\mathrm{Molarity\ (M) = \frac{\text{moles\ of\ solute}}{\text{liters\ of\ solution}}}$

• Example: A 1.00 M NaCl solution contains 1.00 mole of NaCl in 1.00 L of solution.

• To prepare a solution of known molarity, dissolve the calculated mass of solute in enough solvent to reach the desired volume.

Calculating Molarity: Example

• Suppose 24.9 g of Na2SO4 is dissolved in enough water to make 400 mL (0.400 L) of solution.

• Moles of Na2SO4 = $\frac{24.9\ g}{142.04\ g/mol} = 0.175\ mol$

• Molarity = $\frac{0.175\ mol}{0.400\ L} = 0.438\ M$

Using Molarity in Calculations

• Molarity can be used as a conversion factor between moles of solute and volume of solution.

• For example, if a solution is 2.0 M in sugar, 1 L contains 2.0 mol sugar; 0.5 L contains 1.0 mol sugar.

Dilution of Solutions

When a solution is diluted, solvent is added to decrease the concentration of solute. The number of moles of solute remains constant:

$M_1V_1 = M_2V_2$

• $M_1$ and $V_1$ are the initial molarity and volume; $M_2$ and $V_2$ are the final molarity and volume after dilution.

• Example: To dilute 200 mL of 1.50 M NaCl to 400 mL, the final molarity is $M_2 = \frac{(1.50\ M)(0.200\ L)}{0.400\ L} = 0.75\ M$

Solution Stoichiometry

Using Molarity in Chemical Reactions

Molarity allows us to relate the volume of a solution to the number of moles of solute, which can be used in stoichiometric calculations for reactions in solution.

• Example: How many mL of 0.125 M NaOH are needed to react with 24.5 mL of 0.350 M H2SO4?

• Balanced equation: $2\mathrm{NaOH}_{(aq)} + \mathrm{H_2SO_4}_{(aq)} \rightarrow 2\mathrm{H_2O}_{(l)} + \mathrm{Na_2SO_4}_{(aq)}$

• Calculate moles of H2SO4: $0.350\ M \times 0.0245\ L = 0.00858\ mol$

• From stoichiometry: $0.00858\ mol\ H_2SO_4 \times \frac{2\ mol\ NaOH}{1\ mol\ H_2SO_4} = 0.0172\ mol\ NaOH$

• Volume of NaOH needed: $\frac{0.0172\ mol}{0.125\ M} = 0.137\ L = 137\ mL$

Types of Reactions in Aqueous Solution

Precipitation Reactions

• Occur when two solutions of ionic compounds are mixed and an insoluble product (precipitate) forms.

• Example: Mixing solutions of Pb(NO3)2 and KI produces PbI2 (s) as a yellow precipitate.

Predicting Precipitation Reactions

1. Determine the ions present in each reactant.

2. Exchange ions to form possible products.

3. Use solubility rules to determine if any product is insoluble (forms a precipitate).

4. Write the balanced equation, including state symbols: (aq) for soluble, (s) for insoluble.

Writing Chemical Equations

• Molecular equation: Shows all reactants and products as compounds.

• Complete ionic equation: Shows all strong electrolytes as ions.

• Net ionic equation: Shows only the species that actually change during the reaction (removes spectator ions).

Acids and Bases

Definitions

• Arrhenius acid: Produces H+ ions in water (e.g., HCl).

• Arrhenius base: Produces OH- ions in water (e.g., NaOH).

• Brønsted-Lowry acid: Proton donor.

• Brønsted-Lowry base: Proton acceptor.

Strong and Weak Acids/Bases

• Strong acids: Completely ionize in water (e.g., HCl, HNO3, H2SO4).

• Weak acids: Partially ionize in water (e.g., HF, CH3COOH).

• Strong bases: Completely ionize in water (e.g., NaOH, KOH).

• Weak bases: Partially ionize in water (e.g., NH3).

Neutralization Reactions

• When an acid reacts with a base, the products are a salt and water.

• Example: $\mathrm{HCl_{(aq)} + NaOH_{(aq)} \rightarrow NaCl_{(aq)} + H_2O_{(l)}}$

• Net ionic equation: $\mathrm{H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}}$

Gas-Forming Reactions

• Certain salts react with acids to produce gases (e.g., carbonates produce CO2, sulfites produce SO2, sulfides produce H2S).

• Example: $\mathrm{CaCO_3_{(s)} + 2HCl_{(aq)} \rightarrow CaCl_2_{(aq)} + CO_2_{(g)} + H_2O_{(l)}}$

Oxidation-Reduction (Redox) Reactions

Definitions

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Oxidizing agent: Causes oxidation, is itself reduced.

• Reducing agent: Causes reduction, is itself oxidized.

Assigning Oxidation Numbers

• Elements in their standard state: 0

• Monatomic ions: Equal to their charge

• Oxygen: Usually -2 (except in peroxides: -1)

• Hydrogen: +1 (with nonmetals), -1 (with metals)

• Fluorine: -1; other halogens: -1 (unless with oxygen or higher halogen)

• The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion's charge.

Example: Assigning Oxidation Numbers

• In C6H12O6 (glucose): C: ? H: +1 × 12 = +12 O: -2 × 6 = -12 Total: 0, so C = 0

• In SO42-: S: ? O: -2 × 4 = -8 Total: -2, so S = +6

Identifying Redox Changes

• If an atom's oxidation number increases, it is oxidized.

• If an atom's oxidation number decreases, it is reduced.

• Example: $\mathrm{2Na_{(s)} + Cl_2_{(g)} \rightarrow 2NaCl_{(s)}}$ Na: 0 → +1 (oxidized) Cl: 0 → -1 (reduced)