MTTC Chemistry (018) key terms

The digits in a measured quantity that are known with confidence plus one estimated digit; the count signals how precisely an instrument can be read. Leading zeros never count, and trailing zeros count only when a decimal point is present.

How closely a measured result matches the true or accepted value. Poor accuracy usually points to systematic error, such as an uncalibrated balance or a procedure that consistently loses product.

How closely repeated measurements under identical conditions agree with each other. High precision means small random scatter; a data set can be precise yet inaccurate if a systematic offset shifts every reading.

The count of protons in the nucleus of an atom. It defines which element the atom is, and in a neutral atom it also equals the number of electrons surrounding the nucleus.

A form of an element whose atoms have the same proton count but different neutron counts, so their masses differ. Isotopes react almost identically chemically but may differ in nuclear stability.

A mathematical function describing a region of space where an electron is likely to be found. Each orbital holds at most two electrons, which must have opposite spins.

The guideline that electrons occupy the lowest-energy orbitals available before filling higher ones. Applied with the Pauli exclusion principle and Hund's rule, it generates ground-state electron configurations for atoms and ions.

The energy needed to remove the outermost electron from a gaseous atom or ion. Values rise across a period as nuclear charge grows and fall down a group as the outer electron sits farther away.

A relative measure of how strongly an atom pulls shared electrons toward itself within a bond. It increases across a period and decreases down a group; large differences produce ionic character.

The attraction holding oppositely charged ions together in a crystal lattice. It forms when an atom with low ionization energy transfers electrons to an atom with high electron affinity.

A linkage formed when two atoms share one or more pairs of electrons. Each shared pair is attracted to both nuclei, holding the atoms together, and the bond may be polar or nonpolar.

The separation of charge in a covalent bond caused by unequal sharing of electrons between atoms of different electronegativity, giving one end a partial negative charge and the other a partial positive charge.

A model that predicts molecular shape by treating electron pairs around a central atom as domains that repel each other and spread out as far as geometry allows.

The blending of an atom's valence orbitals into a new set of equivalent orbitals with shapes and orientations matching the observed molecular geometry, such as sp3 for tetrahedral carbon.

An unusually strong attraction between a hydrogen atom covalently bound to nitrogen, oxygen, or fluorine and a lone pair on a nearby electronegative atom, responsible for water's high boiling point and surface tension.

The amount of substance equal to Avogadro's number (6.022 × 10^23) of particles. It connects masses measured in the laboratory with counts of atoms, ions, or molecules.

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

The lowest whole-number ratio of the different kinds of atoms in a compound, obtained from percent composition or combustion data rather than from the actual molecular size.

The reactant consumed first when a reaction proceeds, capping how much product can form. Any other reactants remain partly unused and are said to be in excess.

The amount of product predicted from stoichiometry when the limiting reactant is entirely consumed. It is calculated, not measured, and is the denominator in percent-yield calculations.

A signed number assigned to an atom that tracks its electron balance relative to the free element; changes in oxidation state reveal which atoms are oxidized and reduced.

A thermodynamic quantity equal to internal energy plus pressure–volume work. Its change, ΔH, equals the heat exchanged in a process carried out at constant pressure.

The principle that the enthalpy change of an overall reaction equals the sum of the enthalpy changes for any set of steps that add up to that reaction, letting tables of known values be combined.

The minimum energy colliding particles need for a reaction to occur. It sets the height of the barrier between reactants and products and strongly influences how fast the reaction proceeds at a given temperature.

A substance that speeds a reaction by opening a pathway with lower activation energy while emerging unchanged. It shifts neither the equilibrium position nor the equilibrium constant.

The value obtained by multiplying product concentrations, each raised to its coefficient in the balanced equation, and dividing by the analogous reactant term at equilibrium. It varies with temperature and indicates whether products or reactants dominate.

The rule that a system at equilibrium responds to a disturbance in concentration, pressure, or temperature by shifting in the direction that partially counteracts the change.

A thermodynamic function whose change at constant temperature and pressure predicts spontaneity: negative values indicate a spontaneous process, positive values a nonspontaneous one, and zero means equilibrium.

A concentration unit giving the number of moles of solute dissolved per liter of solution. Because solution volume changes with temperature, molarity is slightly temperature dependent.

A logarithmic measure of hydrogen-ion concentration defined as the negative base-10 logarithm of [H+]. At 25 °C values below 7 are acidic, 7 is neutral, and above 7 is basic.

A solution that resists pH change when modest amounts of acid or base are added, usually because it contains a weak acid together with its conjugate base in comparable amounts.

A volumetric procedure in which a standard solution is delivered from a buret to a measured sample until the reaction is just complete, allowing the unknown concentration to be calculated from the volume used.

The equilibrium constant for dissolving a sparingly soluble ionic solid in water. Comparing the reaction quotient with this value predicts whether a precipitate will form.

A solution property that depends on how many solute particles are present, not on what they are; examples include boiling-point elevation, freezing-point depression, vapor-pressure lowering, and osmotic pressure.

A particular arrangement of atoms within an organic molecule that determines its characteristic reactions and physical behavior, such as a hydroxyl, carbonyl, carboxyl, or amino group.

One of two or more compounds that share a molecular formula but differ in how their atoms are connected or arranged in space, giving them distinct properties.

The quality of a molecule that cannot be superimposed on its mirror image, usually because a carbon atom carries four different substituents; the two mirror forms are called enantiomers.

A large molecule built by repeatedly joining small units called monomers through covalent bonds, formed either by addition across double bonds or by condensation with loss of a small molecule.

A biological catalyst, nearly always a protein, that binds specific substrates at an active site and lowers the activation energy of a particular reaction; activity is sensitive to temperature and pH.

A molecule containing both an amino group and a carboxylic acid group; the twenty standard forms link through peptide bonds to build proteins, and the side chain distinguishes one from another.