Solutions and Acids & Bases: Key Concepts and Calculations

Solutions

Solute and Solvent

A solution is a homogeneous mixture composed of two or more substances. The solute is the substance that is dissolved, and the solvent is the substance that does the dissolving (present in greater amount).

• Example: Dissolving 25 g of sugar in 100 g of water: sugar is the solute, water is the solvent.

"Like Dissolves Like"

This phrase means that substances with similar types of intermolecular forces (polarity) will dissolve in each other.

• Polar solutes dissolve in polar solvents (e.g., salt in water).

• Nonpolar solutes dissolve in nonpolar solvents (e.g., oil in hexane).

• Example: Sugar (polar) dissolves in water (polar), but oil (nonpolar) does not.

Solubility of Nonpolar Substances in Water

Nonpolar substances generally do not dissolve in water because water is polar and forms hydrogen bonds, which nonpolar molecules cannot participate in.

• Example: Oil does not dissolve in water.

Electrolytes and Nonelectrolytes

Electrolytes are substances that conduct electricity when dissolved in water because they dissociate into ions. Nonelectrolytes do not conduct electricity because they do not form ions in solution.

• Example: Sodium chloride (NaCl) is a strong electrolyte; sugar is a nonelectrolyte.

Strong and Weak Electrolytes

Strong electrolytes dissociate completely in water, while weak electrolytes only partially dissociate.

• Strong electrolytes: Sodium chloride (NaCl)

• Nonelectrolytes: Sugar (C12H22O11), ethanol (C2H5OH)

Equivalents of Ions

An equivalent is the amount of an ion that supplies 1 mole of electrical charge (+ or -) in solution.

• Example: 1 mole of K+ = 1 equivalent; 1 mole of Ca2+ = 2 equivalents.

Calculating Equivalents in Solution

• To find equivalents: Equivalents = moles of ion × charge

• Example: 1.0 M K2SO4 contains 2 equivalents of K+ per liter.

Charge Balance in Solutions

In any solution, the total equivalents of positive charges must equal the total equivalents of negative charges.

Solubility

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

Saturated Solutions

A saturated solution contains the maximum amount of solute that can dissolve at a given temperature; any additional solute will not dissolve.

• Example: If solubility is 20 g/100 g water, a solution with 19 g is unsaturated, 20 g is saturated, and 25 g is supersaturated (if all dissolves).

Concentration of Solutions

Concentration expresses the amount of solute in a given amount of solution or solvent.

• Common units: mass percent, molarity (M), molality (m), etc.

Calculating Mass Percent

• Formula: $\text{Mass percent} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

• Example: 5 g glucose in 200 g water: $\frac{5}{205} \times 100\% = 2.44\%$

Molarity (M)

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

• Example: 2 moles NaCl in 0.50 L: $M = \frac{2}{0.50} = 4.0$ M

Calculating Moles from Molarity

• Formula: $\text{moles} = M \times \text{volume (L)}$

• Example: 0.80 M solution, 2.25 L: $0.80 \times 2.25 = 1.8$ moles

Dilution Equation

• Formula: $M_1V_1 = M_2V_2$

• Example: 100 mL of 2.0 M HCl diluted to 500 mL: $2.0 \times 0.100 = M_2 \times 0.500$; $M_2 = 0.40$ M

Acids and Bases

Arrhenius Definition

According to Arrhenius, an acid produces H+ ions in water, and a base produces OH- ions.

• Example: HCl → H+ + Cl-; NaOH → Na+ + OH-

Electrolyte Nature of Acids and Bases

Acids and bases are generally electrolytes because they produce ions in solution.

Brønsted-Lowry Definition

In this definition, an acid is a proton (H+) donor, and a base is a proton acceptor.

• Example: NH3 + H2O ⇌ NH4+ + OH-

Conjugate Acid-Base Pairs

Acids and bases exist in pairs that differ by one proton.

• Example: In NH3 + H2O ⇌ NH4+ + OH-, NH3/NH4+ and H2O/OH- are conjugate pairs.

Strong Acids

• Examples: HCl, HNO3, H2SO4 (others: HBr, HI, HClO4)

Electrolyte Strength of Acids

Strong acids are strong electrolytes because they dissociate completely in water.

Auto-dissociation of Water

• Equation: $2H_2O \rightleftharpoons H_3O^+ + OH^-$ or $H_2O \rightleftharpoons H^+ + OH^-$

Ion Product Constant for Water (Kw)

• At 25°C: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$

pH and pOH Calculations

• pH formula: $pH = -\log[H^+]$

• pOH formula: $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$

Acidic, Basic, and Neutral Solutions

• If $[H^+] > 1.0 \times 10^{-7}$ M, solution is acidic (pH < 7).

• If $[H^+] < 1.0 \times 10^{-7}$ M, solution is basic (pH > 7).

• If $[H^+] = 1.0 \times 10^{-7}$ M, solution is neutral (pH = 7).

Sample Calculations

• pH of $[H^+] = 1.0 \times 10^{-9}$ M: $pH = 9$ (basic)

• If pH = 12, $[H^+] = 1.0 \times 10^{-12}$ M

Buffer Solutions

A buffer solution resists changes in pH when small amounts of acid or base are added. Buffers are important in biological systems to maintain stable pH.

• Example: Blood contains a bicarbonate buffer system.

Table: Strong Acids and Their Properties

Acid	Formula	Electrolyte Strength
Hydrochloric acid	HCl	Strong
Nitric acid	HNO3	Strong
Sulfuric acid	H2SO4	Strong
Hydrobromic acid	HBr	Strong
Hydroiodic acid	HI	Strong
Perchloric acid	HClO4	Strong

Additional info: Some explanations and examples were expanded for clarity and completeness based on standard GOB Chemistry curriculum.