Regents Chemistry Ultimate Guide

Dalton’s Theory

Dalton’s Theory

Atoms are indivisible, atoms of an element are identical, compounds are formed from elements, atoms of different elements are different

Thomson’s Theory

Discovered the electron using the Cathode Ray Tube, electron is negative (-) charge and has very little mass

Rutherford’s Theory

Discovered the nucleus through the Gold Foil Experiment. Alpha particles passed through and some were deflected through a sheet of gold, proving the nucleus is small, dense, and positively charged. Atom is mostly empty space otherwise

Bohr Model

Electrons exist in shells around the nucleus. Each energy level has its own amount of energy.

Modern (Wave Mechanical) Model

Orbitals are regions of most probable electron location. An atom consists of a small, positive nucleus surrounded by a cloud of negative electrons.

Average Atomic Mass

Weighted average of all naturally occurring isotopes of an element calculated by (mass * percentage) / 100 + (mass * percentage) / 100.. repeated for as many isotopes

Atomic Spectra

When electrons move from low to high energy they absorb energy, when electrons move from high to low energy levels they release energy

Ground State

The lowest energy state of an atom where electrons fill orbitals following a specific order (2-8-18-32).

Excited State

A higher energy state of an atom where electrons move between shells in a manner that does not adhere to the standard filling order, without changing the total number of electrons.

Elements

Substances that cannot be broken down by chemical means and are found on the periodic table.

Compounds

Substances composed of two or more elements chemically combined, which can be broken down into their constituent elements.

Homogenous

A mixture with no distinguishable difference in appearance or properties.

Heterogeneous

A mixture with differences in appearance and properties.

Solutions

Homogeneous mixtures, such as aqueous solutions, where substances are uniformly distributed.

Protons

Positively charged particles found in the nucleus of an atom.

Neutrons

Electrically neutral particles located in the nucleus of an atom.

Electrons

Negatively charged particles that orbit the nucleus of an atom.

Atom

A neutral particle with no charge

Ions

An atom that has gained or lost electrons and now has a charge

Isotopes

An atom with a different number of neutrons but the same number of protons

Nuclear Charge

The charge of the nucleus, determined by the number of protons (atomic number)

Mass Number

Number of Protons + Neutrons

Atomic Number

Number of protons

Number of neutrons

Mass # - Atomic #

Number of protons

# electrons in a neutral atom

Distillation

Process of separating components based on variations in boiling points.

Filtration

Technique used to separate a solid from a liquid in a heterogeneous mixture.

Chromatography

Method of separating components according to their polarities.

Evaporation/Boiling

Process of separating a homogeneous solution from its solute through vaporization.

Diatomic

Elements that naturally occur as pairs of atoms, not combined in compounds, e.g., Br2, I2, Cl2, H2, O2, F2

Formulas

Represent both qualitative (identity) and quantitative (quantity) information about compounds

Molecular

Indicates the actual number of atoms present in a molecule

Empirical

Represents the simplest whole-number ratio of atoms in a compound

Polyatomic Ions

Found in Table E, should not be separated into individual atoms, contain covalent bonds but can form ionic bonds

Metals

Positive ions that are always written FIRST in chemical formulas.

Non-metals

Elements that are written LAST in chemical formulas.

Criss-Cross Rule

Method of determining subscripts by exchanging the charges of ions in a compound.

Roman Numerals

Used for transition metals with multiple charges to indicate the charge of the metal.

Endothermic Reactions

Reactions where energy is absorbed, with heat written on the left side (reactants), denoted by ΔH = (+).

Exothermic Reactions

Reactions where energy is released, with heat written on the right side (products), denoted by ΔH = (-).

Synthesis

Combination of two substances to form a compound (A + B —> AB)

Decomposition

Compound breaks down into its constituent elements (AB —> A + B)

Single Replacement

Element replaces another in a compound (A + BC —> AB + C)

Double Replacement

Exchange of ions between two compounds (AB + CD —> AD + CB)

Combustion

Reaction of a hydrocarbon with oxygen to produce carbon dioxide and water (Hydrocarbon + O2 —> CO2 + H2O)

Law of Conservation of Mass

States that the mass of all reactants in a chemical reaction must equal the mass of the products.

Solid

Has a definite shape and volume, is rigid, and has fixed patterns.

Liquid

Takes the shape of its container but has a definite volume.

Gas

Has an indefinite shape and volume, and completely fills the container it is in.

Increasing (or decreasing) temperature lines

Kinetic Energy is changing.

Flat lines

Kinetic Energy is constant, but Potential Energy is changing.

Freezing

Transition from liquid to solid, necessitates heat of fusion (334 J)

Melting

Transition from solid to liquid, necessitates heat of fusion (334 J)

Condensation

Transition from gas to liquid, necessitates heat of vaporization (2260 J)

Boiling/Vaporization

Transition from liquid to gas, necessitates heat of vaporization (2260 J)

Sublimation

Transition from solid to gas

Deposition

Transition from gas to solid

Heat of Fusion

The amount of heat required to melt or freeze a substance.

Heat of Vaporization

The amount of heat required to vaporize or condense a substance.

Pressure and Volume

Inverse relationship - as pressure increases, volume decreases, and vice versa.

Pressure and Temperature

Direct relationship - as pressure increases, temperature increases, and vice versa.

Volume and Temperature

Direct relationship - as volume increases, temperature increases, and vice versa.

Alkali Metals

Group 1 elements known for their high reactivity and tendency to form alkaline solutions.

Alkaline Earth Metals

Group 2 elements that are less reactive than alkali metals but still form alkaline solutions.

Transition Metals

Elements in groups 3-12 characterized by their ability to form colored compounds and variable oxidation states.

Halogens

Group 17 elements known for their high reactivity and tendency to form salts with metals.

Nobel Gases

Group 18 elements that are inert, nonreactive, and have a full outer electron shell.

Metals

Malleable, ductile, conductive, luster, low ionization energy and electronegativity, lose electrons and form positive smaller ions

Non-Metals

Brittle, non-conductive, dull, high ionization energy and electronegativity, gain electrons and form negative larger ions

Metalloids

properties of both metals and non-metals

Nobel Gases

unreactive due to full valence shell (8 electrons)

Allotropes

two or more forms of the same element in the same phase, ex. O2 and O3

Breaking Bonds

Absorbs Energy

Forming Bonds

Releases Energy

Covalent

Bond between 2 or more non-metals involving the sharing of electrons.

Ionic

Bond between a metal and a non-metal characterized by the transfer of electrons.

Metallic

Bond found in metals where electrons move freely in a "sea of mobile electrons."

Covalent/Molecular

Low melting point/boiling point, soft, does NOT conduct electricity

Ionic

High melting point/boiling point, hard, conduct electricity only in AQUEOUS or LIQUID phase, NOT solid

Metallic

High melting point/boiling point, hard, conduct in solid and liquid phases

Symmetrical

Non-Polar molecules have a balanced charge distribution

Asymmetrical

Polar molecules have an uneven charge distribution

Polar

Unequal charge distribution with one side more electronegative

Polar Bond

Formed between two different elements

Non-Polar

Equal charge distribution with shared electrons

Non-Polar Bond

Involves the same element on both sides of the bond

Non-Polar Molecule

Exhibits London Dispersion Forces, the weakest intermolecular force

Polar Molecules

Display Dipole-Dipole Forces as intermolecular interactions

Hydrogen Bonding

Present in H2O, NH3, HF molecules where hydrogen is bonded to F, O, or N, known as the strongest intermolecular force

Moles

The amount of a substance that contains Avogadro's number of particles.

Molarity

The concentration of a solution expressed as the number of moles of solute per liter of solution.

Parts Per Million (PPM)

A unit of measurement to represent the concentration of a substance in a mixture, calculated as (part/whole) * 1,000,000 or 1×10^6.

Temperature and Solubility

Increasing temperature generally increases solubility, except for gases.

Gas Solubility

For gases, increasing temperature decreases solubility.

Pressure and Gas Solubility

Increasing pressure enhances the solubility of gases.

"Like Dissolves Like"

Polar solvents dissolve polar solutes, and non-polar solvents dissolve non-polar solutes.

Supersaturated

A solubility reading above the line.

Saturated

A solubility reading on the line.

Unsaturated

A solubility reading below the line.

Concentration

Higher concentration leads to a faster rate of reaction and more effective collisions.