Honors Chemistry EOC Exam Review (Modules 01-11)

Comprehensive vocabulary flashcards covering general chemistry concepts from Modules 01 to 11, including atomic theory, bonding, stoichiometry, gas laws, and solutions.

Density

The amount of matter in a certain volume, calculated as mass divided by volume, and is inversely proportional to volume.

SI Unit for Mass

$$kg$$

SI Unit for Temperature

$$K$$

SI Unit for Volume

$$m^3$$

Direct Relationship

A relationship where both variables increase or decrease together, representing a straight line on a mass vs. volume graph.

Isotopes

Different versions of the same chemical element that contain the same number of protons but a different number of neutrons.

Ion

An atom that has an electrical charge because it has an unequal number of protons and electrons.

Dimitri Mendeleev

The scientist credited with developing the periodic table.

Metals

Elements found in the middle of the periodic table that are shiny, conductive, and form positive ions.

Nonmetals

Elements found on the side of the periodic table that are brittle, dull, nonmalleable, and form negative ions.

Group (or Family)

A vertical column on the periodic table; elements in the same group possess the same number of valence electrons.

Cations

Positively charged ions.

Anions

Negatively charged ions.

Ionic Bond

A bond formed by the transfer of electrons between nonmetals and metals, creating charged ions bound by electrostatic forces.

Molecular (Covalent) Bond

A bond formed when two nonmetals share electrons to create discrete, neutral molecules.

Polyatomic Ion

An ion composed of multiple atoms held together by covalent bonds.

HOFBrINCl

A mnemonic for the seven elements that exist as diatomic molecules: hydrogen, oxygen, fluorine, bromine, iodine, nitrogen, and chlorine.

Metallic Bonding

A lattice of positively charged metal ions surrounded by free-moving valence electrons, resulting in high conductivity, malleability, and luster.

Avogadro’s Number

The number of particles in one mole of a substance.

Molar Mass

The mass in grams of one mole of a substance; for example, the molar mass of $Mg_3(PO_4)_2$ is $262.9\,g\,mol^{-1}$.

Five Indicators of a Chemical Reaction

Temperature change, gas production, precipitate formation, color change, and light production.

Combustion Reaction

A reaction where a hydrocarbon reacts with oxygen to produce carbon dioxide and water, such as $C_5H_{12} + 8O_2 \rightarrow 5CO_2 + 6H_2O$.

Decomposition Reaction

A reaction where a single compound breaks down into two or more products, such as $NiCl_2 \rightarrow Ni + Cl_2$.

Synthesis Reaction

A reaction where two or more substances combine to form a single new substance, such as $2Na + Cl_2 \rightarrow 2NaCl$.

Double Displacement Reaction

A reaction where the ions of two compounds exchange places in an aqueous solution to form two new compounds, such as $Na_2SO_4 + Ca(NO_3)_2 \rightarrow CaSO_4 + 2NaNO_3$.

Single Displacement Reaction

A reaction where one element replaces a similar element in a compound, such as $Na + FeS \rightarrow Na_2S + Fe$.

Law of Conservation of Mass

A principle stating that mass is neither created nor destroyed in a chemical reaction.

Limiting Reactant

The reactant that is completely consumed first in a chemical reaction, limiting the amount of product formed.

Intermolecular Forces

The forces of attraction that hold molecular compounds together.

Dalton’s Law of Partial Pressures

A gas law stating that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual component gases.

Solutions

Homogeneous mixtures composed of solutes dissolved in a solvent.

Molarity

A measure of concentration defined as the number of moles of solute per liter ($dm^3$) of solution.

Dissociation

The process by which ionic compounds separate into their constituent ions when dissolved in water.

Titration

A laboratory technique used to determine the concentration of a solution by reacting it with a standardized solution of known concentration.

Half-life

The time required for half of the atoms of a radioactive isotope to decay.