Regents Chemistry Ultimate Guide

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169 Terms

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Dalton’s Theory

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

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Thomson’s Theory

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

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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

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Bohr Model

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

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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.

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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

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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

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Ground State

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

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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.

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Elements

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

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Compounds

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

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Homogenous

A mixture with no distinguishable difference in appearance or properties.

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Heterogeneous

A mixture with differences in appearance and properties.

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Solutions

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

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Protons

Positively charged particles found in the nucleus of an atom.

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Neutrons

Electrically neutral particles located in the nucleus of an atom.

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Electrons

Negatively charged particles that orbit the nucleus of an atom.

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Atom

A neutral particle with no charge

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Ions

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

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Isotopes

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

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Nuclear Charge

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

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Mass Number

Number of Protons + Neutrons

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Atomic Number

Number of protons

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Number of neutrons

Mass # - Atomic #

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Number of protons

# electrons in a neutral atom

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Distillation

Process of separating components based on variations in boiling points.

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Filtration

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

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Chromatography

Method of separating components according to their polarities.

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Evaporation/Boiling

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

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Diatomic

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

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Formulas

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

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Molecular

Indicates the actual number of atoms present in a molecule

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Empirical

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

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Polyatomic Ions

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

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Metals

Positive ions that are always written FIRST in chemical formulas.

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Non-metals

Elements that are written LAST in chemical formulas.

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Criss-Cross Rule

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

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Roman Numerals

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

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Endothermic Reactions

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

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Exothermic Reactions

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

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Synthesis

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

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Decomposition

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

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Single Replacement

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

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Double Replacement

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

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Combustion

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

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Law of Conservation of Mass

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

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Solid

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

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Liquid

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

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Gas

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

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Increasing (or decreasing) temperature lines

Kinetic Energy is changing.

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Flat lines

Kinetic Energy is constant, but Potential Energy is changing.

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Freezing

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

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Melting

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

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Condensation

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

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Boiling/Vaporization

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

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Sublimation

Transition from solid to gas

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Deposition

Transition from gas to solid

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Heat of Fusion

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

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Heat of Vaporization

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

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Pressure and Volume

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

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Pressure and Temperature

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

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Volume and Temperature

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

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Alkali Metals

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

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Alkaline Earth Metals

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

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Transition Metals

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

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Halogens

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

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Nobel Gases

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

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Metals

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

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Non-Metals

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

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Metalloids

properties of both metals and non-metals

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Nobel Gases

unreactive due to full valence shell (8 electrons)

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Allotropes

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

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Breaking Bonds

Absorbs Energy

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Forming Bonds

Releases Energy

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Covalent

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

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Ionic

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

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Metallic

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

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Covalent/Molecular

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

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Ionic

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

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Metallic

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

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Symmetrical

Non-Polar molecules have a balanced charge distribution

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Asymmetrical

Polar molecules have an uneven charge distribution

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Polar

Unequal charge distribution with one side more electronegative

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Polar Bond

Formed between two different elements

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Non-Polar

Equal charge distribution with shared electrons

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Non-Polar Bond

Involves the same element on both sides of the bond

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Non-Polar Molecule

Exhibits London Dispersion Forces, the weakest intermolecular force

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Polar Molecules

Display Dipole-Dipole Forces as intermolecular interactions

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Hydrogen Bonding

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

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Moles

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

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Molarity

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

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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.

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Temperature and Solubility

Increasing temperature generally increases solubility, except for gases.

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Gas Solubility

For gases, increasing temperature decreases solubility.

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Pressure and Gas Solubility

Increasing pressure enhances the solubility of gases.

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"Like Dissolves Like"

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

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Supersaturated

A solubility reading above the line.

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Saturated

A solubility reading on the line.

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Unsaturated

A solubility reading below the line.

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Concentration

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