Chemistry Physical Setting Regents Exam - Vocabulary Flashcards

Vocabulary flashcards covering key chemical concepts, models, laws, and problem-solving topics from the lecture notes.

Proton

Positively charged subatomic particle located in the nucleus (nucleon).

Neutron

Electrically neutral subatomic particle in the nucleus (nucleon).

Electron

Negatively charged subatomic particle located in orbitals around the nucleus.

Nucleus

Center of an atom that contains protons and neutrons.

Orbitals (clouds)

Regions around the nucleus where electrons are likely to be found.

Mass number

Total number of protons and neutrons in the nucleus.

Atomic number

Number of protons in the nucleus; defines the element.

Neutron number

Number of neutrons in an atom (Mass number minus atomic number).

Isotopes

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

Cation

Positive ion formed when an atom loses electrons; generally smaller than the neutral atom.

Anion

Negative ion formed when an atom gains electrons; generally larger than the neutral atom.

Rutherford’s gold foil experiment

Experiment showing atoms are mostly empty space with a dense, positively charged nucleus.

Thomson’s plum pudding model

Early atomic model with electrons embedded in a positively charged sphere.

Dalton’s atomic model

Solid, uniform sphere of matter as the basic unit of matter.

Bohr model

Model placing electrons in planet-like orbits around the nucleus.

Wave-mechanical model

Current atomic model where electrons occupy orbitals (clouds) around the nucleus.

STP

Standard Temperature and Pressure (273 K, 1 atm).

Bright line spectra

Emission spectra produced when electrons fall to lower energy levels.

Element

Pure substance composed of only one kind of atom.

Binary compound

Compound made of exactly two different kinds of atoms.

Diatomic molecules

Two-atom molecules such as Br2, I2, N2, O2, F2, Cl2, H2.

Significant figures

Rules for determining the number of meaningful digits in a measurement.

Solute

Substance dissolved in a solvent.

Solvent

Substance doing the dissolving (often water).

Isotopic notation (C-14)

Notation showing mass number and atomic number (e.g., 14C with Z=6).

Electron configuration

Arrangement of electrons in energy levels and sublevels.

Mole triangle

Diagram to convert among moles, mass, number of particles, and volume.

Molar volume (22.4 L)

Volume occupied by 1 mole of gas at STP (22.4 L for an ideal gas).

Orbital notation

Depicts electrons as arrows in specific orbitals (1s, 2s, 2p, etc.).

Polyatomic ion

Group of atoms with an overall charge (e.g., NO3-, NH4+).

Coefficient

Number in front of a formula that indicates the ratio of species in a reaction.

Chemical formula neutrality

Formulas arranged so overall charges cancel in ionic compounds.

Binary ionic compound naming

Name the cation first, then the anion with -ide if a binary compound.

Polyatomic ion naming

Keep the polyatomic ion name unchanged when naming compounds.

Physical change

Change in appearance without forming a new substance.

Chemical change

Change that forms new substances.

Reactants and products

Reactants on the left, products on the right of a chemical equation.

Endothermic

Reactions that absorb heat (ΔH positive).

Exothermic

Reactions that release heat (ΔH negative).

Law of Conservation of Mass

Mass of reactants equals mass of products in a chemical reaction.

Gram formula mass

Sum of atomic masses in a formula (g/mole).

Percent composition

Percentage by mass of each element in a compound.

Avogadro’s number

6.02 × 10^23 particles per mole.

States of matter

Solids: definite shape/volume; Liquids: definite volume, no definite shape; Gases: no definite shape or volume.

Solids properties

Rigid, definite shape and volume; particles tightly packed.

Liquids properties

Definite volume but adaptable shape; particles slide past one another.

Gases properties

Widely spaced particles in random motion; easily compressed.

Distillation

Separation technique based on differences in boiling points.

Filtration

Separation of solids from liquids by filtration.

Chromatography

Separation technique based on differential movement through a medium.

Periodic Law

Properties of elements are periodic functions of atomic number.

Periods

Horizontal rows on the Periodic Table.

Groups

Vertical columns on the Periodic Table.

Metals/metalloids/nonmetals

Metals: left of the staircase; Metalloids border the staircase; Nonmetals: above the staircase.

Metallic properties chart

Describes malleability, ductility, luster, conductivity, and ionization energy.

Noble gases

Group 18 elements; inert due to full valence electron shells.

Ionization energy

Energy required to remove an electron; increases up and to the right on the periodic table.

Atomic radii

Size of an atom; generally decreases across a period and increases down a group.

Electronegativity

Tendency of an atom to attract electrons in a bond; increases up and to the right.

Alkali metals

Group 1 metals; highly reactive, soft, form +1 ions.

Alkaline earth metals

Group 2 metals; form +2 ions, reactive but less than Group 1.

Halogens

Group 17 nonmetals; highly reactive, form -1 ions.

Noble gases (Group 18)

Inert gases with complete valence shells.

Valence electrons

Electrons in the outermost shell; last digit of group number approximates their count.

Lewis dot structure

Diagram showing valence electrons as dots around an element symbol.

Kernel

All of an atom except the valence electrons.

Metallic bonds

Bonding in metals described as a lattice of kernels surrounded by a sea of electrons.

Octet rule

Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

Covalent bonds

Bond formed by sharing a pair of electrons between atoms.

Ionic bonds

Bond formed by transfer of electrons from a metal to a nonmetal.

Nonpolar covalent bonds

Covalent bonds formed between identical atoms or with very small electronegativity difference.

Polar covalent bonds

Covalent bonds with unequal sharing due to electronegativity difference (roughly 0.4–1.7).

Molecular substances

Substances held together mainly by covalent bonds.

Ionic compounds

Compounds composed of cations and anions bonded ionically.

Hydrogen bonds

Strong dipole-dipole attractions when H is bonded to N, O, or F; high boiling/melting points.

Solubility and like dissolves like

Substances tend to dissolve in solvents with similar properties.

Solubility with temperature

For most solids, solubility increases with higher temperature.

Solubility of gases

Solubility of gases generally increases with higher pressure and lowers with higher temperature.

Molarity

Concentration = moles of solute per liter of solution.

Percent by mass

Mass of solute divided by total mass of solution, times 100%.

ppm

Parts per million; a mass-based concentration unit.

Boiling point elevation / freezing point depression

Colligative properties where solutes raise bp and lower mp of solvents.

Normal boiling point

The boiling point of a substance at 1 atm pressure.

Reaction rate factors

Concentration, surface area, temperature, pressure, and catalysts can affect rate.

Activation energy

Minimum energy required for a reaction to occur.

Equilibrium

Rates of forward and reverse reactions are equal; concentrations remain constant.

Le Chatelier’s principle

If a system at dynamic equilibrium is disturbed, it shifts to counteract the change.

Keq (equilibrium constant)

Keq = [C]^y[D]^z / [A]^w[B]^x for a given reaction wA + xB ⇄ yC + zD.

Oxidation

Loss of electrons; oxidation number increases.

Reduction

Gain of electrons; oxidation number decreases.

LEO says GER

Losing Electrons = Oxidation; Gaining Electrons = Reduction.

Oxidizing vs reducing agents

Oxidizing agents are reduced; reducing agents are oxidized in a redox reaction.

Electrochemical cells

Cells that produce electricity from spontaneous redox reactions; include anode and cathode.

Anode / Cathode

Anode is where oxidation occurs; cathode is where reduction occurs.

Electrolysis

Using electrical energy to drive a nonspontaneous redox reaction.

Acids and bases (Arrhenius)

Acids produce H+ in solution; bases produce OH− in solution.

Acids and bases (Brønsted)

Acids donate protons; bases accept protons.

Neutralization

Acid-base reaction forming water and a salt.

Titrations

Controlled neutralization to determine the concentration of an acid or base.

Organic carbon-containing compounds

All organic compounds contain carbon and typically hydrogen.