Comprehensive Chemistry: Solutions, Electrolytes, and Acid-Base Reactions

Solution

a homogeneous mixture where a solute is uniformly mixed with a solvent.

Solute

substance present in a smaller amount in a solution (ex: sugar in sweet tea).

Solvent

substance present in a larger amount in a solution (ex: water in sweet tea).

Aqueous solution

a solution in which water is the solvent.

Colloid

homogeneous mixture with particles 1-1000 nm in diameter (ex: milk).

Suspension

mixture with large particles that will settle over time (ex: blood).

Solvation

process where solvent particles surround and interact with solute particles.

Hydration

solvation that occurs when water is the solvent.

Ion-dipole interaction

attraction between a charged ion and polar water molecules.

Like dissolves like

polar substances dissolve in polar solvents; nonpolar in nonpolar.

Unsaturated solution

contains less than the maximum solute that can dissolve.

Saturated solution

contains the maximum solute that can dissolve at that temperature.

Temperature effect on solubility

↑ T → ↑ solubility for most solids, ↓ solubility for gases.

Pressure effect on solubility

↑ pressure → ↑ gas solubility (Henry's Law).

Henry's Law

the solubility of a gas in a liquid is directly proportional to the gas's pressure above the liquid.

(applied) Soda

In soda, CO₂ stays dissolved under pressure; when opened, pressure drops and gas escapes.

Electrolyte

substance that forms ions in water and conducts electricity.

Strong electrolyte

completely dissociates into ions (ex: NaCl → Na⁺ + Cl⁻).

Weak electrolyte

partially dissociates; equilibrium between ions and molecules.

Nonelectrolyte

dissolves in water but does not produce ions (ex: glucose).

Equivalent (Eq)

relates ion charge to number of moles of ions.

mEq/L

milliequivalents per liter; used in IV fluids and blood analysis.

(applied) Ion equivalents

1 mol Mg²⁺ = 2 Eq Mg²⁺; 1 mol Na⁺ = 1 Eq Na⁺.

Mass percent (m/m)

(mass solute / mass solution) × 100%.

Volume percent (v/v)

(volume solute / volume solution) × 100%.

Mass/volume percent (m/v)

(grams solute / mL solution) × 100%.

Molarity (M)

moles solute / liters solution.

ppm/ppb

parts per million or billion; used for trace concentrations.

(applied) Glucose concentration

5 % (m/v) glucose = 5 g glucose / 100 mL solution.

Dilution

adding solvent to reduce concentration.

Dilution formula

C₁V₁ = C₂V₂.

(applied) IV solutions

used to make IV solutions or lab dilutions from a stock solution.

Isotonic solution

same solute concentration as inside the cell; no net water movement.

Hypotonic solution

lower solute concentration outside the cell; water moves in → cell swells.

Hypertonic solution

higher solute concentration outside the cell; water leaves → cell shrinks (crenates).

Osmosis

movement of water through a semipermeable membrane from low to high solute concentration.

Diffusion

movement of solute from high to low concentration (passive).

Active transport

movement against a concentration gradient using energy (ATP).

Dialysis

movement of solutes and water through a membrane (medical use: removes waste from blood).

Arrhenius acid

produces H⁺ (= H₃O⁺) ions in water.

Arrhenius base

produces OH⁻ ions in water.

Bronsted-Lowry acid

proton donor.

Bronsted-Lowry base

proton acceptor.

Amphoteric

can act as acid or base (ex: water).

Strong acid/base

completely dissociates in water.

Weak acid/base

partially dissociates (reversible equilibrium).

Strong acids

HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.

Strong bases

LiOH, NaOH, KOH, RbOH, CsOH.

Neutralization reaction

acid + base → salt + water.

Reaction rate

speed at which reactants form products.

Factors affecting rate

temperature, reactant concentration, catalyst.

Reversible reaction

proceeds in both forward and reverse directions.

Chemical equilibrium

rate forward = rate reverse; concentrations constant.

Equilibrium constant (K)

ratio (products/reactants) at equilibrium.

Le Chatelier's Principle

if equilibrium is disturbed, system shifts to restore balance.

Water autoionization

H₂O ⇌ H₃O⁺ + OH⁻.

pH

−log[H₃O⁺]; measures acidity.

Ka (acid dissociation constant)

ratio of products/reactants for weak acid.

pKa

−log Ka; smaller pKa = stronger acid.

Conjugate acid-base pair

differ by one proton (H⁺).

Buffer solution

resists changes in pH when small acid or base is added.