Chemistry Regents Vocab (copy)

Elements

Pure substances composed of only one type of atom and cannot be broken down.

Diatomic Elements

Elements that form two atom molecules in their natural form at STP. Remember the phrase “BrINClHOF”.

Binary Compounds

Substances made up of two kinds of atoms that have chemically combined.

Solutions/Homogeneous Mixtures

Made up of different substances but there are no visible differences, meaning its uniform throughout.

Heterogeneous Mixtures

Visibly have different components and are not uniform throughout.

Distillation

Separates mixtures in which the components have different boiling points.

Filtration

Separates mixtures of solids and liquids (different sizes): therefor, filtration cannot be used to separate solutions.

The Law of Conservation

The mass of the reactants in a chemical reaction is always equal to the mass of the products. Energy and charge are also conserved.

Physical Changes

Do not form new substances (ex. phase changes, dissolving)

Chemical Changes

Result in the formation of new substances (ex. burning, reacting)

Solids

1. Have particles that are very close together in a regular, geometric pattern

2. Have a definite shape and definite volume

Liquids

1. Have closely-spaces particles that can easily slide past one another

2. Have indefinite shape, but have a definite volume

Gasses

1. Have widely-spaced particles that are in random motion and completely fill a container

2. Easily compressed and have indefinite shape and indefinite volume

Sublimation

When a substance turns from a solid directly into a gas

Deposition

When a substance turns from a gas directly into a solid

Heat Flow

Heat energy flows from substances of a high temperature to substances of a low temperature

Kinetic Energy (KE)

1. Related to temperature

2. On the heating/cooling curves- on the slopes KE changes while on the flat lines (phase changes) KE is constant.

Potential Energy (PE)

1. Related to the distance between particles

2. On the heating/cooling curves- on the slopes PE is constant while on the flat lines (phase changes) PE changes

The Kinetic Molecular Theory

Ideal gases move in constant, random, straight-line motion, have elastic collisions (no energy is lost), have negligible (very small) volume, no attractive forces, and collide with container walls causing pressure.

Real Gasses

1. Do not behave like ideal gases because they have some volume and some attraction

2. Behave more like ideal gasses at low pressures and high temps (PLIGHT)

Avogadro’s Hypothesis

Equal volume of gases at the same temp and pressure contain an equal number of molecules

Relation of Pressure, Temp, and Volume

1. As pressure on a gas increases, volume decreases (inverse relation)

2. As the pressure on a gas increases, temperature increases (direct relation)

3. As the temperature of a gas increases, volume increases (direct relation)

Boiling

1. Liquids boil when their vapor pressure is equal to the atmospheric pressure

2. The normal bp of a substance is the temp at which it boils at 1 atm

3. The stronger attraction between molecules, the higher its mp and bp

Dalton’s Model

1. The First Model

2. Solid sphere of matter that was uniform throughout

Thompsons Model

1. The Second Model

2. Discovered the electron and developed the “plum-pudding” model

Ruthford’s Gold Foil Experiment

1. The Third Model

2. Showed that an atom is mostly empty space with a small, dense, positively charged nucleus

The Bohr Model

1. The Fourth Model

2. Placed electrons in “planet-like”orbitals around the nucleus

The Wave-Mechanical Model

1. The Most Recent Model

2. Has electrons in “clouds” (orbitals) around the nucleus

Protons

1. Positively changed and have a mass of 1 amu

2. Atomic #

3. In an atom’s nucleus (nucleons)

4. In a neutral atom = to # of electrons

5. Nuclear charge is = to # of protons

Neutrons

1. Have no charge (neutral) and have a mass of 1 amu

2. In the atom’s nucleus (nucleons)

3. # of neutrons= mass #- atomic #

Electrons

1. Negatively charged and have a mass of 0 amu

2. Found in “clouds”(orbitals) around the nucleus

3. In a neutral atom = to # of protons

Mass Number

Protons+Neutrons

Nuclear Charge

# of protons

Cations

1. Positive ions and form when a neutral atom loses electrons.

2. Smaller than the parent atom.

3. To calculate # of electrons- subtract charge from atomic #

Anions

1. Negative ions that form when a neutral atom gains electrons

2. Larger than the parent atom

3. To calculate # of electrons- add charge to atomic #

Isotopes

Atoms with the same # of protons (atomic #) but a different number of neutrons (mass #)

Average Atomic Mass

The weighted average of the atomic masses of all naturally occurring isotopes.

The Bright Line Spectrum

Electrons emit energy as they move from higher energy levels (excited state) back down to lower (ground state) energy levels.

Alpha Particles

Have a mass of 4 amu and a charge of 2+

Beta Particles

Have a mass of 0 amu and a charge of -1.

Posistrons

Have a mass of 0 and and a charge of 1+

Transmutation

When one element changes into another

1. Neutral Transmutation- happens on its own

2. Artificial Transmutation- is when a nucleus is bombarded with particles to force the decay

Half-Life

The length of time it takes for one half of the mass of a sample to radioactively decay.

Nuclear Reactions

Produce more energy than chemical reactions because a small amount of mass is converted into a large amount of energy.

Fission Reactions

Split heavy nuclei into smaller ones

Fusion Reactions

Occur when light nuclei combine to form a heavier nucleus

Groups on Period Table

Same # valence electrons so similar properties

1- Alkali Metals

2- Alkaline Earth Metals

3-12- Transition Metals (colored compounds ans solutions)

17- Halogens

18- Noble Gasses (nonreactive and stable)

Properties of Metals and Non-Metals

Allotropes

Forms of the same element in the same phase with different structures and different properties

Atomic Radius

Size of an atom. Increases down a group and decreases across a period.

Electronegativity

Measure of an element’s attraction for electrons. The higher the electronegativity is the stronger attraction for electrons. It decreases down a group and increases across a period.

Ionization Energy

The amount of energy needed to remove the outermost electron. It decreases down a group and increases across a period.

Energy

Absorbed when a bond breaks- substance is less stable

Releases when bond is formed- substance becomes more stable

BARF

Ionic Bonds

When metals transfer electrons to nonmetals

Ionic Compounds

Have an END of 1.7 or greater

Covalent Bonds

Form when two atoms share a pair of electrons. Substances containing mostly covalent bonds are called molcular.

Nonpolar Covalent Bonds

When two of the same nonmetals bond together. The electronegativity difference is 0 and electrons are shared equally.

Polar Covalent Bonds

When different nonmetals bond together. The END is between 0.4 and 1.7 and electrons are shared unequally. The atom with the higher EN gets a slight negative charge and the atom with the lower EN gets a slight positive charge, This molecule is called a dipole.

Non Polar/ Polar Molecules

Non-Polar is symmetrical

Polar is asymmetrical

SNAP

Ionic vs. Covalent Properties

Polyatomic Ions

Groups of atoms with a net charge that act as one unit. Any compound that contains a polyatomic ion contains both ionic and covalent bonds.

Metallic Bonds

Consist of metal ions surrounded by a “sea” of mobile valence electrons.

Van der Waals Forces

1. Dipole-Dipole Forces

   1. Occur between polar molecules

2. Dispersion Forces

   1. Temporary forces that occur between nonpolar molecules. The larger the non polar molecules are, the stronger the dispersion forces between them.

Hydrogen Bonds

Forms when a hydrogen atom of one molecule is attracted to N,O, or F in another molecule. This gives the compound unusually high melting and boiling points.

Molecule-ion attraction

Occur between a polar liquid and an ionic substance. This attraction is responsible for dissolving and is therefore present in all solutions (aq).

Reactants

Left side of the reaction arrow

Products

Right side of reaction arrow

Synthesis Reactions

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

Decomposition Reactions

Occur when a single reactant forms two or more products

Single Replacement Reactions

Occurs when one element replaces another in a compound. The free element needs to be above the element it is trying to replace on Table J.

Double Replacement Reaction

Occurs when two compounds react to form two new compounds. One of the products must be water, a gas, or a precipitate (insoluble on Table F).

Gram Formula Mass

The sum of the atomic masses of all the atoms in a substance. Each atomic number needs to be multiplied by the number of atoms present.

Hydrates and Anhydrous Substances

Compounds with water inside. Anhydrous substances have no water inside and are formed by heating a hydrate. To find the mass of water, subtract th mass of the anhydrate from the mass of the hydrate.

Empirical vs Molecular Formulas

In empirical formulas all subscripts are reduced. In molecular formulas the subscripts represent the actual ratio of atoms.

Avogadro’s Number

6.02×10²³. One mole of any substance occupies 22.4 L

Solute vs Solvent

The solute is the substance being dissolved while the solvent is the substance that dissolved the solute.

Like Dissolves Like

Polar solvents dissolve polar/ ionic solutes and nonpolar solvents dissolve nonpolar solutes.

Temp Affects Solubility

Temp increase- solubility increase for solids and liquids but decreases for gases.

Gases are most soluble under low temps and high pressure

Concentration

A measure of the amount of solute in a solution. Concentrated solutions contain a large amount of solute, while dilute solutions contain a small amount of solute.

Increasing the Concentration of a Solution

Decreases the fp and increases the bp. The greater concentration, the greater the effect on the freezing and boiling point. This is called a colligative property.

Arrhenius Acids

Produce H⁺ (hydrogen) ions in a solution, also called H30⁺ (hydronium ions)

Arrhenius Bases

Produce OH⁻ (hydroxide) ions in a solution

Properties of Acids

1. Taste Sour

2. Turn Litmus Red

3. Turn Phenolphthalein Colorless

4. pH less than 7 (lower pH=more H⁺ ions)

5. Produce H⁺ and H3O⁺ ions in solution

6. Donate H⁺ (alt theory)

7. Electrolytes (conduct electricity)

Properties of Bases

1. Taste Bitter

2. Turn Litmus Blue

3. Turn Phenolphthalein Pink

4. pH greater than 7 (higher pH=more OH⁻ ions)

5. Produce OH⁻ ions in solution

6. Accept H⁺ (alt theory)

7. Electrolytes (conduct electricity)

8. Don’t react with metals

Titration

A lab technique used to find the concentration of an acid or base sample.

Collision Theory

For a reaction to occur, particles must collide with enough energy and at proper angles

Catalysts

Speed up reactions by lowering activation energy and creating an alternate pathway.

Have no effect on equilibrium; speed up forward and reverse reactions equally.

Reaction Rate

1. Increase concentration= increase reaction rate

2. Increase surface area= increase reaction rate

3. Increase temp= increase reaction rate

Endothermic Reactions

Absorb heat and the energy value is added to the left (reactant) side of the reaction arrow in a forward reaction.

Exothermic Reaction

Release energy and the energy value is added to the right (product) side of the reaction arrow in a forward reaction.

The Heat of Reaction

The difference between the potential energy of the products and the potential energy of the reactants. For endothermic reactions it is positive and for exothermic reactions it is negative.

Entropy

The measure of randomness or disorder. Natural tendency towards greater disorder. Gas > liquid> solid also aq solution > solute&solvent

Spontaneous Reactions

Favor low energy and high entropy

Equilibrium

Reaction rates are equal at equilibrium while concentrations are constant.

Adding any reactant or product to a system at equilibrium will shift the equilibrium away from the added substance. (Add Away)

Removing any reactant or product from a system at equilibrium will shift the equilibrium toward that removed substance. (Take Towards)

Increase in temp shifts a system in the endothermic direction (away from heat)

Decrease in temp shifts a system in the exothermic direction (toward heat)

Increase the pressure on a gaseous equilibrium will shift the equilibrium toward the side with fewer moles of gas. (volume was decreased.

Decrease the press on a gaseous equilibrium will shift the equilibrium toward the side with more moles of gas. (volume was increased)

Redox Reactions

Involves the exchange of electrons (Lose Electrons Oxidation, Gain Electrons Reduction). The number of electrons lost = the number of electrons gained.

Oxidation

The loss of electrons (LEO). The oxidation number decincreasesreases as a result. The electrons are added to the right side of the reaction arrow.

Reduction

The gain of electrons (GER). The oxidation number decreases as a result. The electrons are added to the left side of the reaction arrow.

Oxidizing Agents and Reducing Agents

Oxidizing Agents- Get reduced

Reducing Agents- Get oxidized

Voltaic Cells

Convert chemical to electrical energy with a spontaneous redox reaction.

The anode (negative) is the higher metal on Table J, it gets oxidized

The cathode (positive) is the lower metal on Table J, it gets reduced

(ANode OXidation LEO, BIG REDuction CAThode GER)

Electrons flow through a wire from the anode to the cathode

The salt bridge allows for the migration of neutral ions

Electrolytic Cells

Use an applied electrical current (battery) to force a nonspontanious redox reaction to take place (converts electrical to chemical energy).

Used for metal plating

Anode (positive)- used for plating

Cathode (negative)- object that gets coated