ILTS Science: Chemistry (240) flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Independent variable

The factor the experimenter deliberately changes or manipulates to test its effect on the outcome; all other conditions are held constant to isolate cause and effect.

Accuracy versus precision

Accuracy is closeness of a measurement to the accepted value; precision is closeness of repeated measurements to one another. Systematic error harms accuracy, random error harms precision.

Scientific theory

A well-substantiated explanation of natural phenomena supported by extensive, reproducible evidence. Theories explain why patterns occur and are not upgraded into laws.

Significant figures rule

All nonzero digits and interior zeros count; leading zeros do not; trailing zeros count only with a decimal point. Multiplication keeps the least significant figures; addition keeps the least decimal places.

Percent error

Absolute difference between experimental and accepted values divided by the accepted value, multiplied by 100. It quantifies measurement accuracy relative to a true standard.

Isotopes

Atoms of the same element with identical proton number but different neutron numbers, giving different mass numbers. Their weighted abundances determine the element's average atomic mass.

Aufbau principle

Electrons fill orbitals in order of increasing energy, from lowest to highest. Combined with Hund's rule and the Pauli exclusion principle, it predicts ground-state electron configurations.

Effective nuclear charge

The net positive charge an outer electron experiences after inner electrons shield the nucleus. It increases across a period, shrinking atomic radius and raising ionization energy.

Mole concept

One mole contains 6.022 × 10²³ representative particles. Molar mass converts between grams and moles, enabling stoichiometric calculations from balanced equations.

VSEPR theory

Electron domains around a central atom repel and arrange to maximize separation, setting molecular geometry. Lone pairs occupy more space than bonding pairs, compressing bond angles.

Limiting reactant

The reactant consumed first, determined by comparing moles of product each reactant can form. It caps the theoretical yield; other reactants remain in excess.

Net ionic equation

An equation showing only the species that actually react, with spectator ions removed. It must balance both atoms and charge and reveals precipitate or gas formation.

Hydrogen bonding

A strong dipole attraction between a hydrogen bonded to N, O, or F and a lone pair on a nearby electronegative atom. It raises boiling points and surface tension.

Molarity

Moles of solute divided by liters of solution, expressed in mol/L. Dilution follows M₁V₁ = M₂V₂ because solute moles stay constant when solvent is added.

Enthalpy change

Heat exchanged at constant pressure; negative for exothermic reactions and positive for endothermic ones. Hess's law sums stepwise changes to find the overall value.

Gibbs free energy

ΔG = ΔH − TΔS determines spontaneity at constant temperature and pressure. A negative ΔG means the process is spontaneous; ΔG° = −RT ln K links it to equilibrium.

Collision theory

Reactions occur when particles collide with sufficient energy to overcome activation energy and with correct orientation. Catalysts lower activation energy and speed the rate.

Le Chatelier's principle

A system at equilibrium shifts to counteract a disturbance such as concentration, pressure, or temperature change. Heating favors the endothermic direction; pressure favors fewer gas moles.

Reaction quotient Q

The equilibrium expression evaluated at current concentrations. Comparing Q to K predicts direction: Q < K shifts forward, Q > K shifts reverse, Q = K means equilibrium.

Colligative properties

Properties depending on solute particle count, not identity: boiling-point elevation, freezing-point depression, vapor-pressure lowering, and osmosis. Formulas include van't Hoff factor i.

Buffer system

A weak acid with its conjugate base (or weak base with conjugate acid) that resists pH change. Henderson–Hasselbalch gives pH = pKa + log([A⁻]/[HA]).

pH scale

pH = −log[H₃O⁺]; at 25 °C pH + pOH = 14 and Kw = 1.0 × 10⁻¹⁴. Values below 7 are acidic, above 7 basic.

Electrochemical cell

Oxidation occurs at the anode and reduction at the cathode. E°cell = E°cathode − E°anode; a positive value indicates a spontaneous galvanic cell with ΔG° = −nFE°cell.

Titration equivalence point

The point where acid and base have reacted in exact stoichiometric proportions. For a strong acid–strong base titration at 25 °C, the equivalence pH is 7.

Functional group

A specific atom grouping in an organic molecule that determines characteristic reactivity. Examples include hydroxyl, carbonyl, carboxyl, amino, ester, and amide groups.

Peptide bond

A covalent amide linkage formed between the carboxyl group of one amino acid and the amino group of another, releasing water. It builds polypeptide chains in proteins.

Enzyme catalysis

Enzymes are specific protein catalysts that lower activation energy by stabilizing the transition state. Denaturation disrupts tertiary structure and destroys activity.

Alpha decay

Emission of a helium-4 nucleus, reducing mass number by 4 and atomic number by 2. It changes the element identity and is common in heavy unstable nuclei.

Half-life

The time for half of a radioactive sample to decay, following first-order kinetics. After n half-lives, the fraction remaining is (1/2)^n.

Acid dilution safety

Always add acid slowly to water, never water to acid. The highly exothermic hydration can boil and spatter if concentrated acid contacts water directly.

Distillation

A separation technique exploiting differences in boiling point. The more volatile component vaporizes first and is condensed, separating it from less volatile substances.

5E instructional model

A sequence of Engage, Explore, Explain, Elaborate, and Evaluate. Exploration with evidence precedes formal explanation, supporting inquiry-based chemistry learning.