Regents Chemistry Comprehensive Review

Flashcards covering essential vocabulary and concepts from Units 1-15 for the Chemistry Regents exam.

Element

A substance composed of only one type of atom that cannot be broken down by chemical means.

Compound

A substance made of two or more elements chemically bonded in a fixed ratio that can be broken down by chemical change.

Mixture

Two or more substances physically mixed in no fixed ratio that retain individual properties and are separable by physical means.

Homogeneous Mixture

A mixture that is uniform throughout, also known as a solution.

Distillation

A physical separation technique that separates liquids based on different boiling points.

Filtration

A process used to separate a solid from a liquid using a filter.

Chromatography

A method of separating mixtures based on solubility and polarity.

Allotropes

Different forms of the same element in the same phase, such as $O_2$ and $O_3$ (ozone) or diamond and graphite.

Physical Property

A characteristic observed without changing the identity of the substance, such as color, density, melting point, or hardness.

Chemical Property

A characteristic that describes how a substance reacts with something else, such as flammability or reactivity with acid.

Chemical Change

A change that results in the formation of a new substance, indicated by clues like color change, gas bubbles, or heat release.

Kinetic Molecular Theory (Solids)

Particles are in fixed positions (crystal lattice), only vibrate, and have definite shape and volume.

Temperature

A measure of the average kinetic energy of the particles in a substance.

Sublimation

An endothermic phase change where a solid turns directly into a gas, such as dry ice ($CO_2$).

Deposition

An exothermic phase change where a gas turns directly into a solid.

Heating Curve (Flat Parts)

Sections of a phase change diagram where temperature/Average Kinetic Energy stays constant while Potential Energy changes during a phase change.

Density

A characteristic property calculated using the formula $D = \frac{m}{V}$, expressed in units like $g/mL$ or $g/cm^3$.

Proton

A subatomic particle with a charge of $+1$ and a mass of $1\text{ amu}$, located in the nucleus, which determines the atomic number.

Neutron

A subatomic particle with no charge and a mass of $1\text{ amu}$, located in the nucleus.

Electron

A subatomic particle with a charge of $-1$ and negligible mass ($\text{~}0\text{ amu}$), located in orbitals around the nucleus.

Mass Number

The sum of the protons and neutrons in an atom's nucleus.

Isotopes

Atoms of the same element with the same atomic number but a different number of neutrons (and thus different mass numbers).

Thomson (Plum Pudding Model)

An atomic model that discovered the electron, depicting the atom as a positive sphere with negative electrons embedded.

Rutherford (Gold Foil Experiment)

An experiment showing atoms are mostly empty space with a tiny, dense, positive nucleus.

Wave-Mechanical Model

The modern cloud model where electrons are found in orbitals, which are the most probable locations of electrons.

Valence Electrons

Electrons found in the outermost shell or energy level of an atom.

Excited State

A state where one or more electrons have absorbed energy and jumped to a higher energy level, resulting in a configuration that does not match the Periodic Table.

Bright-line Spectrum

A unique fingerprint of light emitted by an element when its excited electrons fall back to the ground state.

Cation

A positive ion formed when a metal atom loses electrons.

Anion

A negative ion formed when a nonmetal atom gains electrons.

Alpha Particle

A type of nuclear decay with a mass of $4$ and charge of $+2$ ($^{4}_{2}He$); it has the lowest penetrating power.

Gamma Radiation (\boldsymbol{\text{\gamma}})

Pure energy (electromagnetic radiation) with zero mass and zero charge; it has the highest penetrating power.

Natural Transmutation

A process where one radioactive reactant spontaneously decays into two products.

Half-Life

The time required for half of a radioactive sample to decay, values for which are found on Table N.

Fission

A nuclear reaction where a heavy nucleus splits into smaller nuclei, neutrons, and energy.

Fusion

A nuclear reaction where light nuclei combine to form a heavier nucleus, producing more energy than fission.

Periods

Horizontal rows (1-7) on the Periodic Table indicating the number of electron shells.

Groups

Vertical columns (1-18) on the Periodic Table; elements in the same group share the same number of valence electrons and similar chemical properties.

Metalloids

Elements located on the staircase (B, Si, Ge, As, Sb, Te) that have properties of both metals and nonmetals.

Atomic Radius

A periodic trend that decreases left to right across a period and increases down a group.

Ionization Energy

The energy required to remove an electron from an atom; it increases left to right and decreases down a group.

Electronegativity

The attraction for electrons in a bond; Fluorine has the highest value at $4.0$.

Ionic Bond

A bond formed by the transfer of electrons from a metal to a nonmetal or polyatomic ion.

Covalent Bond

A bond formed by the sharing of electrons between two nonmetals.

Metallic Bond

A bond characterized by a "sea of mobile electrons" around positive metal ions, allowing for conductivity and malleability.

SNAP Rule

A mnemonic for molecular polarity: Symmetrical = Nonpolar, Asymmetrical = Polar.

Hydrogen Bonding

The strongest intermolecular force, occurring when hydrogen is bonded to Fluorine, Oxygen, or Nitrogen ($F$, $O$, or $N$).

Synthesis

A chemical reaction type where two or more reactants combine to form one product ($A + B \rightarrow AB$).

Decomposition

A chemical reaction type where one compound breaks down into two or more products ($AB \rightarrow A + B$).

Mole

A counting unit equal to $6.02 \times 10^{23}$ particles; one mole of a gas at STP occupies $22.4\text{ L}$.

Gram Formula Mass (GFM)

The sum of the atomic masses of all atoms in a chemical formula, expressed in $g/mol$.

Empirical Formula

The simplest whole-number ratio of atoms in a compound.

Limiting Reactant

The reactant that runs out first in a chemical reaction and determines the amount of product formed.

Ideal Gas (PLIGHT)

An imaginary gas that follows KMT perfectly; favored under conditions of Pressure Low, Ideal (no IMF), and High Temperature.

Avogadro's Hypothesis

The principle that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Solubility (Like Dissolves Like)

The rule that polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.

Molarity (M)

A measure of concentration defined as the moles of solute per liter of solution.

Colligative Properties

Properties such as Boiling Point Elevation and Freezing Point Depression that change when a solute is added to a solvent.

Catalyst

A substance that speeds up a reaction by providing an alternative pathway with a lower activation energy, without being consumed.

Exothermic Reaction

A reaction where energy is released (-\boldsymbol{\text{\Delta}} H), and the potential energy of the products is lower than the reactants.

Endothermic Reaction

A reaction where energy is absorbed (+\boldsymbol{\text{\Delta}} H), and the potential energy of the products is higher than the reactants.

BARF Rule

Bonds Absorb energy when Breaking; Release energy when Forming.

Entropy

A measure of disorder or randomness; gases have the highest and solids have the lowest.

Dynamic Equilibrium

A state in a closed system where the rates of the forward and reverse reactions are equal and concentrations remain constant.

Le Chatelier's Principle

The rule that if stress is applied to a system at equilibrium, the system shifts to relieve that stress.

Specific Heat Capacity ($C$)

The amount of heat needed to raise the temperature of $1\text{ gram}$ of a substance by $1\text{ degree Celsius}$; for water, it is 4.18\text{ J/g}\text{\cdot}\text{C}.

Arrhenius Acid

A substance that produces $H^+$ (or $H_3O^+$) ions as the only positive ions in aqueous solution.

Electrolytes

Substances (acids, bases, and salts) that conduct electricity when dissolved in aqueous solution.

Neutralization

The reaction of an acid and a base to produce water and a salt ($\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$).

Titration

A laboratory technique used to determine the concentration of an unknown acid or base, using the formula $M_aV_a = M_bV_b$.

Oxidation

The loss of electrons resulting in an increase in oxidation number (LEO).

Reduction

The gain of electrons resulting in a decrease in oxidation number (GER).

Voltaic Cell

An electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction.

Electrolytic Cell

An electrochemical cell that uses electrical energy to drive a non-spontaneous chemical reaction, often used for electroplating.

Hydrocarbons

Organic compounds containing only carbon and hydrogen.

Saturated Hydrocarbon

An organic compound where all carbon-carbon bonds are single bonds (Alkanes).

Isomers

Compounds with the same molecular formula but different structural arrangements, leading to different properties.

Saponification

An organic reaction specifically for making soap from fat (ester) and a base.

Polymerization

A reaction where many small molecules (monomers) join together to form a long chain (polymer).