MTEL Chemistry key terms

The amount of a substance containing 6.022 × 10^23 representative particles, such as atoms or molecules. It links particle counts to measurable mass and is the basis for stoichiometric calculation.

The mass in grams of one mole of a substance, numerically equal to its formula mass in atomic mass units. It converts between grams and moles in calculations.

The value 6.022 × 10^23, the number of particles in one mole of a substance. It allows conversion between particle count and amount in moles.

The simplest whole-number ratio of atoms of each element in a compound. It is found by converting percentage composition to moles and dividing by the smallest value.

The actual number of atoms of each element in one molecule of a compound. It is a whole-number multiple of the empirical formula, determined using molar mass.

The mass percentage of each element in a compound, calculated by dividing the total mass of that element in one mole by the molar mass. It helps identify unknown compounds.

The quantitative relationship among reactants and products in a balanced chemical equation. Mole ratios from coefficients allow prediction of reactant needs and product yields.

The reactant that is completely consumed first and therefore determines the maximum amount of product formed. It is identified by comparing available mole ratios to those in the equation.

The actual yield of a reaction divided by the theoretical yield, multiplied by 100. It measures reaction efficiency and reflects losses from side reactions or incomplete reaction.

A concentration unit equal to moles of solute per liter of solution. It is commonly used in solution stoichiometry and is temperature dependent because volume changes with temperature.

Adding solvent to a solution to lower its concentration without changing the amount of solute. The relationship M1V1 = M2V2 describes the process.

A procedure that determines the concentration of an unknown solution by adding a standard solution until the reaction reaches its equivalence point, often signaled by an indicator.

The point in a titration where the moles of titrant added exactly equal the moles needed to react with the analyte according to the balanced equation. It may differ from the endpoint.

The equation PV = nRT relating pressure, volume, moles, and absolute temperature of a gas. The constant R is 0.0821 L·atm/(mol·K).

The total pressure of a gas mixture equals the sum of the partial pressures of its components. It is used with gas collection over water and mixtures of nonreacting gases.

The rate of effusion of a gas is inversely proportional to the square root of its molar mass. Lighter gases effuse faster than heavier ones at the same temperature.

The absolute temperature scale whose zero is absolute zero, with no negative values. Gas law calculations require temperatures expressed in kelvin, converted by adding 273 to Celsius.

The number of protons in the nucleus of an atom, which identifies the element. In a neutral atom it also equals the number of electrons.

The sum of protons and neutrons in an atom's nucleus. Isotopes of an element differ in mass number while sharing the same atomic number.

An atom of the same element with a different number of neutrons, and therefore a different mass number. Isotopes share chemical properties but may differ in stability.

The energy required to remove an electron from a gaseous atom or ion. It generally increases across a period and decreases down a group.

A measure of the size of an atom, typically half the distance between nuclei of bonded atoms. It decreases across a period and increases down a group.

The tendency of an atom in a bond to attract shared electrons. Differences in electronegativity determine whether a bond is nonpolar covalent, polar covalent, or ionic.

The electrostatic attraction between oppositely charged ions formed by transfer of electrons, usually between a metal and a nonmetal. Such compounds form crystal lattices with high melting points.

A covalent bond in which electrons are shared unequally because the atoms differ in electronegativity, creating partial charges. It gives molecules a dipole moment if not canceled by geometry.

A measure of the separation of positive and negative charge in a molecule, arising from bond dipoles and molecular geometry. Symmetric molecules may have polar bonds yet no net dipole.

A model predicting molecular shape by arranging electron domains around a central atom as far apart as possible. Lone pairs occupy more space than bonding pairs and alter bond angles.

The mixing of atomic orbitals to form new equivalent orbitals that explain molecular geometry, such as sp, sp2, and sp3. It accounts for bond angles and multiple bonding.

The energy released when gaseous ions combine to form a solid ionic lattice, or the energy needed to separate them. It increases with higher ion charges and smaller ionic radii.

Attractions between molecules, including London dispersion forces, dipole–dipole interactions, and hydrogen bonding. They are weaker than covalent bonds but determine boiling point and solubility.

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

The pressure exerted by a vapor in equilibrium with its liquid at a given temperature. It increases with temperature and is lower for liquids with stronger intermolecular forces.

The temperature at which a liquid's vapor pressure equals the external pressure. It rises with stronger intermolecular forces and with greater external pressure.

When a system at equilibrium is disturbed by changes in concentration, pressure, or temperature, it shifts in the direction that partially counteracts the disturbance. It predicts equilibrium responses.

The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its coefficient power. Its value depends on temperature and indicates reaction extent.

A decrease in solubility or a shift in equilibrium caused by adding an ion already present in the system. It follows from Le Châtelier's principle.

A logarithmic measure of hydrogen ion concentration equal to the negative log of [H+]. Values below 7 are acidic and above 7 are basic at standard conditions.

A solution that resists large pH changes when small amounts of acid or base are added, usually containing a weak acid and its conjugate base. Its capacity depends on component concentrations.

The negative logarithm of the acid dissociation constant Ka, indicating acid strength. A lower pKa means a stronger acid; pKa helps predict buffer regions.

The loss of electrons by a species, accompanied by an increase in oxidation state. It always occurs together with reduction in a redox reaction.

The gain of electrons by a species, accompanied by a decrease in oxidation state. In a redox reaction it occurs simultaneously with oxidation.

A number assigned to an atom that reflects its electron count relative to the neutral element, following set rules. It is used to identify redox changes and balance equations.