203 Ways to Pass the Chemistry Regents Exam Flashcards

Vocabulary-style practice flashcards based on the Chemistry Regents Exam review notes.

Nucleons

Protons and neutrons located in an atom’s nucleus.

Electrons

Small, negatively charged ($-$) particles found in "clouds" (orbitals) around an atom’s nucleus.

Orbitals

The "clouds" around an atom’s nucleus where electrons are found according to the current wave-mechanical model.

Mass number

The sum of an atom’s number of protons and neutrons.

Atomic number

A value equal to the number of protons in the nucleus of an atom.

Number of neutrons

Calculated by subtracting the atomic number from the mass number.

Isotopes

Atoms with equal numbers of protons but different numbers of neutrons.

Cations

Positive ($+$) ions formed when a neutral atom loses electrons; they are smaller than their parent atom.

Anions

Negative ions formed when a neutral atom gains electrons; they are larger than their parent atom.

Ernest Rutherford’s gold foil experiment

An experiment showing that an atom is mostly empty space with a small, dense, positively-charged nucleus.

J.J. Thompson

The scientist who discovered the electron and developed the "plum-pudding" model of the atom.

Dalton’s model

An atomic model described as a solid sphere of matter that was uniform throughout.

Bohr Model

An atomic model that placed electrons in "planet-like" orbits around the nucleus.

STP

Standard Temperature and Pressure, defined as $273 Kelvin$ and $1 atm$.

Bright line spectra

Light produced when electrons emit energy as they jump from higher energy levels back down to lower energy levels.

Binary compounds

Substances made up of only two kinds of atoms, such as $H_2O$, $NH_3$, and $CO_2$.

Diatomic molecules

Elements that form two-atom molecules in their natural form at STP: $Br_2, I_2, N_2, Cl_2, H_2, O_2, F_2$ (BrINClHOF).

Pacific Atlantic Rule

A rule for counting significant figures: start from the Pacific (left) if a decimal is present, or the Atlantic (right) if it is absent.

Homogeneous mixtures

Mixtures that are uniform throughout, with solutions like air or salt water being the best examples.

Heterogeneous mixtures

Mixtures with discernable components that are not uniform throughout, such as soil or vegetable soup.

Solute

The substance being dissolved in a solution.

Solvent

The substance that dissolves the solute in a solution.

Electron configuration

The distribution of electrons in an atom, found at the bottom center of an element’s box on the periodic table.

Polyatomic ions

Groups of atoms with an overall charge, listed in Table E, such as $NO_3^{1-}$ or $SO_4^{2-}$.

Coefficients

Numbers written in front of formulas in chemical equations to give the ratios of reactants and products.

Physical changes

Changes that do not form new substances and merely change the appearance of the original material, like melting ice.

Chemical changes

Changes that result in the formation of new substances, such as burning hydrogen gas.

Endothermic reactions

Reactions that absorb heat, where the energy value is on the left side of the reaction arrow.

Exothermic reactions

Reactions that release energy, where energy is listed as a product in the reaction.

Synthesis reactions

Reactions occurring when two or more reactants combine to form a single product.

Decomposition reactions

Reactions occurring when a single reactant forms two or more products.

Law of Conservation of Mass

The principle stating that the mass of reactants is always equal to the mass of products in a chemical equation.

Gram formula mass

The sum of the atomic masses of all atoms in a substance, measured in $g/mole$.

Avogadro’s number

$6.02 \times 10^{23}$, representing the number of particles in $1 mole$ of a substance.

Sublimation

The phase change where a substance turns from a solid directly into a gas, characteristic of $CO_2$ and $I_2$.

Specific heat capacity formula

$q = mc\text{Δ}t$, used to calculate heat absorbed or released by substances.

Heat of fusion

The heat absorbed/released when $1 gram$ of a substance changes between the solid and liquid phases, which is $334 J/g$ for water.

Heat of vaporization

The heat absorbed/released when $1 gram$ of a substance changes between the liquid and gaseous phases, which is $2260 J/g$ for water.

Combined gas law

The formula $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$, where temperature must always be in Kelvins.

Distillation

A method to separate mixtures with different boiling points.

Filtration

A method to separate mixtures of solids and liquids.

Periodic Law

The law stating that the properties of elements are periodic functions of their atomic numbers.

Noble gases

Group 18 elements that are inert and stable because their valence-level electrons are completely filled.

Electronegativity

A measure of an element’s attraction for electrons, which increases as you move up and to the right on the Periodic Table.

Covalent bonds

Bonds formed when two atoms share a pair of electrons.

Ionic bonds

Bonds formed when one atom transfers an electron to another, typically with an electronegativity difference greater than $1.7$.

Molarity

A measurement of solution concentration equal to the number of moles of solute divided by the liters of solution.

Catalysts

Substances that speed up reactions by lowering their activation energies and can be reused many times over.

Entropy ($S$)

A property that is high in unorganized systems, such as a gas.

Oxidation

The loss of electrons by an atom or ion, resulting in an increase in the oxidation number.