Solutions Flashcards
(17 cards)
Definition of a Solution
A solution is a homogeneous mixture in which the components are dispersed at the molecular or ionic scale. Its composition is uniform throughout, so any sample contains the same proportions of solute and solvent. Light passes without scattering if the particles are truly molecularly dispersed.
Molecular Events During Solution Formation
When a solution forms, solvent particles surround and separate individual solute particles, dispersing them randomly through the medium. Attractive forces between solute and solvent replace or exceed those within the pure substances, making the mixture energetically favourable. This molecular mingling is called solvation—or hydration when water is the solvent.
Solute and Solvent
The solute is the component present in lesser amount and being dissolved, whereas the solvent is the component in greater amount that does the dissolving. In aqueous solutions water is always the solvent regardless of proportion. Their roles are defined functionally rather than strictly by phase.
Saturated vs Unsaturated Solutions
A saturated solution contains the maximum solute that can dissolve under the given conditions, so dissolution and crystallisation occur at equal rates. An unsaturated solution holds less solute than that limit and can dissolve more. Supersaturated solutions exceed saturation temporarily but are unstable.
K_sp and Solubility
A low K_sp value shows that only a tiny amount of solid must dissolve before the ionic product equals K_sp and precipitation becomes favourable, so the compound is poorly soluble. A high K_sp means greater solubility because ions can accumulate to higher concentrations before equilibrium is reached. Comparing the ionic reaction quotient Q to K_sp predicts whether precipitation will occur.
Arrhenius Acids and Bases
In the Arrhenius model, an acid is a substance that increases the concentration of hydrogen ions (H⁺) in aqueous solution, whereas a base increases the concentration of hydroxide ions (OH⁻). These definitions are restricted to water and to species that directly release those ions. They do not adequately describe non‑hydroxide bases such as ammonia.
Brønsted‑Lowry Acids and Bases
The Brønsted‑Lowry definition broadens the concept: an acid is a proton (H⁺) donor and a base is a proton acceptor, irrespective of solvent. Acid‑base reactions are therefore proton‑transfer processes. This framework explains the basicity of substances that bind protons without supplying hydroxide.
Acid or Base Added to Water
Adding an acid to water produces hydronium ions (H₃O⁺); the extent depends on the acid’s strength. Strong acids dissociate nearly completely, giving a sharp pH drop, whereas weak acids establish an equilibrium that limits hydronium concentration. Adding a base produces hydroxide ions, raising pH according to the base’s strength.
Acid Added to Base
Mixing an acid with a base results in neutralisation: H⁺ from the acid combines with OH⁻ from the base to form water, while the remaining ions form a salt. For a strong acid‑strong base pair this reaction proceeds essentially to completion, bringing the solution toward neutral pH. If either partner is weak, the final pH depends on the relative strengths and concentrations involved.
Strong vs Weak Acid or Base
A strong acid or base dissociates essentially 100 % in water, so almost every formula unit splits into ions. A weak one ionizes only partially, establishing an equilibrium between ions and intact molecules. The pH or pOH therefore changes far more for a strong reagent at the same concentration because the ionic species responsible for acidity or basicity are more abundant.
Self‑Ionization of Water and K₍w₎
Even pure water undergoes a tiny equilibrium where two H₂O molecules trade a proton, forming hydronium (H₃O⁺) and hydroxide (OH⁻). The equilibrium constant for this process is K₍w₎=[H₃O⁺][OH⁻]=1.0 × 10⁻¹⁴ at 25 °C. Because both ions appear in equal amounts, each has a concentration of 1.0 × 10⁻⁷ M in neutral water.
Definition of pH
pH is the negative base‑10 logarithm of the hydronium‑ion concentration, pH=−log[H₃O⁺]. The scale is inverse, so larger [H₃O⁺] means smaller pH; neutral water at 25 °C sits at pH 7. Each integer step reflects a ten‑fold change in acidity, making the scale compact and easy to compare.
Three Laws of Thermodynamics
First: energy is conserved; the internal energy change of a system equals heat plus work transferred. Second: every spontaneous process increases the entropy of the universe, imposing a direction on natural change. Third: as a perfect crystal approaches absolute zero, its entropy trends toward zero, providing an absolute entropy reference.
Concentrations from K₍eq₎
Set up an ICE (Initial–Change–Equilibrium) table with algebraic changes, insert the expressions into the K₍eq₎ formula, and solve for the unknown x. The quadratic may be simplified if K₍eq₎ is very small or large. The resulting equilibrium concentrations complete the table.
Ion Concentrations from K₍sp₎
Again assume the molar solubility is s and express each ion’s concentration in terms of s, substitute into the K₍sp₎ expression, and solve for s. Multiplying s by the appropriate coefficient gives individual ion concentrations. Common‑ion situations require adding any pre‑existing ion concentration before solving.
Concentration from Dissociation Constant and pH
Use pH to find [H⁺] and set that equal to x in the ICE table for HA dissociation. Plug x and the given K₍a₎ into the equilibrium expression, solve for the initial concentration or remaining amount as required. The algebra links observed acidity to how much acid was present.
Components of a Voltaic Cell
Given a diagram, locate the anode (labelled negative, oxidation occurs), cathode (positive, reduction), external circuit for electrons, salt bridge or porous cup for ion migration, and half‑cell solutions. The metal that loses mass is the anode; the one gaining mass is the cathode. Direction of electron flow always anode → cathode.
