Pharmacist Licensure Examination Reviewer - VOCABULARY Flashcards (General, Inorganic, Organic, and Medicinal Chemistry Concepts)

A comprehensive set of vocabulary-style flashcards covering key topics from general chemistry, inorganic/organic chemistry, and medicinal chemistry as presented in the lecture notes. Each card defines a term and explains its relevance to PhLE exam topics.

Matter

Anything that has mass and occupies space; the basic subject of chemistry.

Atom

The basic unit of an element that retains the properties of that element; composed of protons, neutrons, and electrons.

Element

A pure substance made of one kind of atom that cannot be broken down by simple chemical means.

Compound

A substance formed by chemically joining two or more elements in fixed proportions.

Mixture

Matter composed of two or more substances not chemically bonded; can be homogeneous or heterogeneous.

Homogeneous Mixture

A uniform mixture with one phase (e.g., solution).

Heterogeneous Mixture

A mixture with two or more distinct phases.

Law of Definite Proportions

Elements in a compound combine in fixed whole-number ratios.

Law of Multiple Proportions

Elements can combine in different ratios to form different compounds.

Law of Conservation of Mass

In a chemical reaction, mass is conserved; total mass of reactants equals total mass of products.

Physical Change

A change in matter that alters form but not chemical identity (e.g., phase change).

Chemical Change

A process that forms new substances with new properties; evidenced by gas evolution, precipitate, heat, or color change.

Phase/State

Solid, liquid, gas; plasmas and, in some contexts, liquid crystals as a mesophase.

Colloid

A heterogeneous mixture with particles dispersed in a medium; larger than in solutions but smaller than in suspensions.

Tyndall Effect

Light scattering by colloidal particles, giving a visible beam in a colloid.

Brownian Movement

Random zigzag motion of colloidal particles due to collisions with molecules.

Electrophoresis

Movement of charged particles through a medium under an electric field.

SDS-PAGE

A gel electrophoresis technique used to separate proteins and nucleic acids by size.

Covalent Bond

A bond formed by the sharing of electron pairs between atoms.

Ionic Bond

A bond formed by electrostatic attraction between oppositely charged ions.

Hydrogen Bond

A strong dipole-dipole interaction involving hydrogen and a highly electronegative atom (N, O, F).

Van der Waals Forces

Weak intermolecular forces including London dispersion, dipole-dipole, and dipole-induced dipole attractions.

Dipole Moment

A measure of the separation of positive and negative charges in a molecule.

Electronegativity

Tendency of an atom to attract electrons in a chemical bond.

Ionization Energy

Energy required to remove an electron from a neutral atom in the gaseous state.

Electron Affinity

Energy change when an electron is added to a neutral atom to form an anion.

Periodic Table

A tabular arrangement of elements organized by increasing atomic number and periodic properties.

Group/Family

Column of the periodic table; elements share similar valence electron configurations.

Period

Row of the periodic table; elements show a progression of properties across the row.

Isotope

Atoms of the same element with the same number of protons but different numbers of neutrons.

Isobar

Different elements with the same atomic mass number but different atomic numbers.

Isotone

Different elements with the same number of neutrons.

Electron Configuration

Distribution of electrons among atomic orbitals in an atom.

Quantum Numbers

Numbers describing electron position: n (principal), l (azimuthal), ml (magnetic), ms (spin).

Pauli Exclusion Principle

No two electrons in an atom can have the same set of quantum numbers; each orbital holds at most 2 electrons.

Aufbau Principle

Electrons fill lowest energy levels first, building up to higher levels.

Hund’s Rule

Orbitals of the same subshell are singly occupied before pairing occurs.

Hund’s Rule of Maximum Multiplicity

Electrons fill degenerate orbitals singly before pairing.

Hybridization

Mixing of atomic orbitals (e.g., sp3, sp2, sp) to form new hybrid orbitals.

Sigma Bond (σ)

A single covalent bond formed by end-to-end overlap along the bond axis.

Pi Bond (π)

A covalent bond formed by side-by-side overlap of p orbitals; accompanies a σ bond in multiple bonds.

Molecular Orbital

A region in a molecule where electrons are most likely to be found; formed by overlapping atomic orbitals.

Valence Electrons

Electrons in the outermost shell that determine chemical reactivity.

Reaction Rate

Speed at which a chemical reaction proceeds.

Collision Theory

Reaction rate depends on the frequency and energy of molecular collisions.

Activation Energy (Ea)

Minimum energy required for reactants to transform into products.

Transition State

High-energy, unstable arrangement of atoms at the top of the reaction energy barrier.

Rate Law

Expression of the rate of reaction as a function of reactant concentrations; Rate = K[A]^x[B]^y.

Catalyst

Substance that speeds a reaction by providing an alternative pathway with lower Ea.

Le Chatelier’s Principle

A system at equilibrium responds to stresses (concentration, temperature, pressure) by shifting to counteract the stress.

Thermodynamics

Science of energy transformations and how energy affects matter and its properties.

System vs Surroundings

System is the part being studied; surroundings include everything else.

State Function

Path-independent properties (e.g., H, S, G) that depend only on the current state.

Non-State Function

Path-dependent quantities (e.g., work, heat) that depend on the process taken.

First Law of Thermodynamics

Energy cannot be created or destroyed; it can only be transformed.

Second Law of Thermodynamics

Spontaneous processes increase the total entropy of the universe.

Third Law of Thermodynamics

The entropy of a perfect crystal at absolute zero is zero.

Enthalpy (H)

Heat content of a system at constant pressure.

Entropy (S)

A measure of the degree of disorder or randomness of a system.

Gibbs Free Energy (G)

Energy that determines spontaneity; ΔG = ΔH − TΔS; ΔG = 0 at equilibrium.

Internal Energy (U)

Total energy contained within a system (thermal + chemical energy).

Heat (q)

Energy transferred due to temperature difference; positive if absorbed, negative if released.

Work (w)

Energy transfer due to pressure-volume work; sign depends on whether work is done by or on the system.

pH

A measure of hydrogen ion concentration; pH = −log[H+].

pOH

A measure of hydroxide ion concentration; pH + pOH = 14 at 25°C.

Buffer Capacity

Ability of a buffer to resist changes in pH upon addition of acid or base.

Henderson–Hasselbalch Equation

pH = pKa + log([salt]/[acid]); relates pH to buffer components.

Amphoteric Substance

A substance that can act as either an acid or a base.

Isohydric Solution

A solution with the same pH as a standard solution.

Solubility

Maximum amount of solute that can dissolve in a solvent at a given temperature.

Miscibility

Ability of two substances to mix in all proportions; form a homogeneous solution.

Raoult’s Law

Lowering of solvent vapor pressure is proportional to the mole fraction of solute.

Henry’s Law

Solubility of a gas in a liquid is proportional to the partial pressure of the gas.

Vapor Pressure Lowering

Decrease in solvent vapor pressure due to nonvolatile solute, a colligative property.

Boiling Point Elevation

Increase in boiling point caused by added solute (colligative property).

Freezing Point Depression

Lowering of freezing point due to solute presence (colligative property).

Osmotic Pressure

Pressure needed to prevent solvent flow across a semipermeable membrane.

Ideal Gas Law

PV = nRT; relates pressure, volume, temperature, and moles of gas.

Avogadro’s Law

Volume is directly proportional to the number of gas moles.

Boyle’s Law

Volume is inversely proportional to pressure (at fixed temperature and amount of gas).

Charles’s Law

Volume is directly proportional to temperature (in Kelvin) at constant pressure.

Grahams Law

Rate of effusion or diffusion is inversely proportional to the square root of molar mass.

Synthesis (Combination) Reaction

Direct union of elements to form a compound.

Decomposition Reaction

Complex substance breaks down into simpler substances.

Single Replacement Reaction

An element replaces another in a compound (A + BC → B + AC).

Double Replacement Reaction

Exchange of partners between two compounds (AB + CD → AD + CB).

Neutralization

Acid reacts with base to form a salt and water.

Redox (Oxidation-Reduction)

Reactions involving electron transfer; oxidation is loss of electrons, reduction is gain.

Lucas Test

Qualitative test for alcohols using ZnCl2 in HCl to differentiate 1°, 2°, 3° alcohols.

Alcohols

Hydroxyl-containing organic compounds (−OH); classified as primary, secondary, or tertiary.

Aldehydes & Ketones

Functional groups with carbonyl C=O; aldehydes have –CHO, ketones have C=O within the chain.

Carboxylic Acids

R–COOH; acidic organic acids; derivatives include esters, amides, anhydrides.

Esters

R–COOR; formed by reaction of carboxylic acids with alcohols; often have pleasant odors.

Amides

Derivatives of carboxylic acids with –CONR2; formed by condensation with amines.

Amines

Organic derivatives of ammonia (RNH2, R2NH, R3N); basic in nature.

IUPAC Nomenclature

Systematic naming of organic compounds using longest carbon chain, lowest locants, and suffix changes for functional groups.

Isomerism

Compounds with same molecular formula but different structures; includes structural, geometric, and stereoisomers.

Stereochemistry

Study of spatial arrangement of atoms; includes enantiomers, diastereomers, and epimers.

Optical Activity

Rotation of polarized light by chiral molecules; measured with a polarimeter.

Chirality/Chiral Center

A carbon atom bonded to four different groups; can have non-superimposable mirror images (enantiomers).