BSSE Chemistry 1429 key terms

The number of protons in the nucleus of an atom, which uniquely identifies the element. In a neutral atom, it equals the number of electrons. It determines the element's position in the periodic table.

The total number of protons and neutrons in an atomic nucleus. It is always a whole number and is used to distinguish isotopes of the same element. It is approximately equal to the atomic mass in atomic mass units.

Atoms of the same element that have the same number of protons but different numbers of neutrons. They have identical chemical properties but different masses. Some isotopes are radioactive and are used in dating and medicine.

The arrangement of electrons in an atom's orbitals, described by quantum numbers and typically written using noble gas notation. It follows the Aufbau principle, Hund's rule, and the Pauli exclusion principle. It determines chemical properties and reactivity.

A region of space around the nucleus where an electron is most likely to be found. Each orbital can hold a maximum of two electrons with opposite spins. Orbitals are labeled s, p, d, f and have characteristic shapes.

A regular variation in properties such as atomic radius, ionization energy, and electronegativity across a period or down a group in the periodic table. These trends arise from changes in effective nuclear charge and electron shielding.

A measure of an atom's ability to attract shared electrons in a chemical bond. It generally increases across a period and decreases down a group. Differences in electronegativity determine bond polarity.

The electrostatic attraction between oppositely charged ions formed by the transfer of electrons from a metal to a nonmetal. It typically results in a crystalline solid with high melting point. It is non-directional.

A chemical bond formed by the sharing of one or more pairs of electrons between atoms. It is directional and typically occurs between nonmetals. The shared electrons occupy molecular orbitals.

Valence Shell Electron Pair Repulsion theory predicts the geometry of molecules by assuming that electron pairs around a central atom repel each other and arrange themselves as far apart as possible. It accounts for lone pair effects on bond angles.

The mixing of atomic orbitals on a central atom to form new equivalent hybrid orbitals that are used for bonding. It explains observed molecular geometries such as tetrahedral, trigonal planar, and linear. Common types are sp, sp2, and sp3.

The attractive forces between molecules, including London dispersion forces, dipole-dipole interactions, and hydrogen bonding. They are weaker than covalent bonds but determine physical properties such as boiling point, melting point, and solubility.

A strong dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom. It significantly raises boiling points and affects structures like DNA and water.

The equation PV = nRT that describes the behavior of an ideal gas, where P is pressure, V is volume, n is moles, R is the gas constant, and T is temperature in kelvin. It combines Boyle's, Charles's, and Avogadro's laws.

The amount of substance containing Avogadro's number (6.022 × 10^23) of particles. It is the base unit for amount in the SI system and links mass to number of particles. It is central to stoichiometric calculations.

The mass of one mole of a substance, expressed in grams per mole. It is numerically equal to the atomic or molecular mass in atomic mass units. It is used to convert between mass and moles in stoichiometry.

The quantitative relationship between reactants and products in a chemical reaction, based on the balanced chemical equation. It allows calculation of amounts of substances consumed or produced. It is essential for predicting yields and limiting reactants.

The reactant that is completely consumed in a chemical reaction and thus determines the maximum amount of product that can be formed. Other reactants are in excess. Identifying it is key to calculating theoretical yield.

The heat content of a system at constant pressure, denoted by H. The change in enthalpy (ΔH) equals the heat absorbed or released in a process. It is used to describe exothermic and endothermic reactions.

The principle that the total enthalpy change of a reaction is independent of the pathway, depending only on the initial and final states. It allows calculation of ΔH for reactions that are difficult to measure directly by combining known enthalpy changes.

A chemical reaction that releases heat to the surroundings, resulting in a negative enthalpy change (ΔH < 0). The products have lower enthalpy than the reactants. Combustion is a common example.

A dynamic state in a reversible reaction where the rates of the forward and reverse reactions are equal, so the concentrations of reactants and products remain constant. The equilibrium constant expression quantifies the position of equilibrium.

When a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the system shifts to counteract the disturbance and re-establish equilibrium. It predicts the direction of shift but not the rate.

A logarithmic scale measuring the acidity of a solution, defined as pH = -log[H+]. A pH below 7 is acidic, 7 is neutral, and above 7 is basic. It is crucial in acid-base chemistry and many biological processes.

A solution that resists changes in pH when small amounts of acid or base are added, typically containing a weak acid and its conjugate base (or a weak base and its conjugate acid). It works by neutralizing added H+ or OH-.

An equation that relates the reaction rate to the concentrations of reactants, with each concentration raised to a power determined experimentally. The overall order is the sum of these exponents. It cannot be deduced from the stoichiometry alone.

The minimum energy required for a chemical reaction to occur, representing the energy barrier between reactants and the transition state. It is denoted by Ea and affects the reaction rate according to the Arrhenius equation.

A substance that increases the rate of a reaction by providing an alternative pathway with lower activation energy, without being consumed in the overall reaction. It does not change the equilibrium position or the thermodynamics of the reaction.

The process in which an atom, ion, or molecule loses electrons, resulting in an increase in oxidation state. It occurs at the anode in an electrochemical cell. It is always accompanied by reduction.

The process in which an atom, ion, or molecule gains electrons, resulting in a decrease in oxidation state. It occurs at the cathode in an electrochemical cell. It is always accompanied by oxidation.

A number assigned to an element in a compound that represents the number of electrons lost or gained by an atom of that element. It is used to track electron transfer in redox reactions. Rules exist for assigning oxidation states.

The potential of a half-cell measured relative to the standard hydrogen electrode under standard conditions (1 M, 1 atm, 25 °C). It indicates the tendency of a species to be reduced. It is used to predict the direction of redox reactions.

A specific group of atoms within a molecule that is responsible for the characteristic chemical reactions of that molecule. Examples include hydroxyl, carbonyl, carboxyl, and amino groups. It determines the molecule's properties and reactivity.

Molecules that have the same molecular formula but different arrangements of atoms, leading to different structures and properties. Types include structural isomers, geometric isomers, and optical isomers. They are important in organic chemistry and drug design.

A systematic set of rules for naming chemical compounds based on their structure, developed by the International Union of Pure and Applied Chemistry. It ensures unambiguous communication of chemical identity. It involves identifying the longest chain and principal functional group.

A laboratory technique for determining the concentration of a solution by reacting it with a standard solution of known concentration. The equivalence point is often detected using an indicator or a pH meter. It is widely used in acid-base and redox analysis.

The digits in a measured value that carry meaningful information about its precision, including all certain digits plus one uncertain digit. They are used to report results with the correct level of confidence. Rules govern their use in calculations.

The degree to which repeated measurements under the same conditions give similar results, reflecting the reproducibility of a measurement. It is often expressed as standard deviation. It is distinct from accuracy.