Topic 1: Atomic Structure and Matter

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, which proves that 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 it’s 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:

(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

Electron Configuration

Ground State (lowest energy): 2-8-18-32

Excited State (higher energy): when electrons move around shells and don’t follow the order above. THE NUMBER OF ELECTRONS DO NOT CHANGE.

Matter

Elements: Cannot be broken down by chemical means (on the periodic table)

Compounds: Composed of 2 or more elements chemically combined; can be broken down into elements.

Both of these are pure substances.

Mixtures

Combination of 2 or more pure substances than can be physically separated

Homogenous: no distinguishable difference in appearance/properties

Heterogenous: difference in appearance and properties

All solutions (aqueous) are homogenous mixtures.

Particles

Protons are positive (in nucleus)

Neutrons are neutral (in nucleus)

Electrons are negative (outside nucleus)

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 different number of neutrons but same number of protons

Nuclear Charge - Charge of nucleus, only protons = atomic number

Mass Number - (top number) Number of Protons + Neutrons

Atomic Number - (bottom number) Number of protons

Number of neutrons = Mass # - Atomic #

Number of protons = # electrons in a neutral atom

Physical Separation Methods

Distillation: separation of components by differences in boiling points

Filtration: Separating a solid from a liquid in a heterogenous mixture

Chromatography: separation of components by polarities

Evaporation/Boiling: separating a homogenous solution from it’s solute

Topic 2: Nomenclature and Formula Writing

Diatomic - Occur in nature this way, not combined in a compound: Br₂, I₂, Cl₂, H₂, O₂, F₂

Formulas - Show qualitive info (what it is) and quantitative info (how many there are)

Molecular - How many atoms are actually there

Empirical - A simplified molecular formula

Polyatomic Ions - Listed in Table E, Do NOT break up, contains covalent bonds and forms ionic bonds

Writing Formulas

Metals (positive ions) are always written FIRST, Non-metals are are written LAST

Criss-Cross the charges of ions and make them into subscripts

Mg⁺² + N^-3 —> Mg₃N₂

Roman Numerals

Written when the metal has more than one charge (transition metals). Roman numerals state the charge of the metal. DO NOT USE if the metal has only 1 charge

Fe⁺² + O^-2 —> Iron (II) Oxide

Fe⁺³ + O^-2 —> Iron (III) Oxide

Endothermic Reactions - Energy is absorbed; energy (heat) is written on the left (reactants).

ΔH = (+)

Exothermic Reactions - Energy is released; energy (heat) is written on the right side (products).

ΔH = (-)

Types of Chemical Reactions

Synthesis: A + B —> AB

Decomposition: AB —> A + B

Single Replacement: A + BC —> AB + C (element + compound —> compound + element)

Double Replacement: AB + CD —> AD + CB (compound + compound —> compound + compound)

Combustion: Hydrocarbon + O₂ —> CO₂ + H₂O

All Single Replacement reactions are Redox Reactions

All Neutralization Reactions are Double Replacement Reactions

Law of Conservation of Mass - Mass of all reactants must Equal Mass of Products

Balancing Equations

Place coefficients in front of compounds or elements. Do NOT change subscripts. Coefficient distributes into the entire compound for each atom. Balance atoms on each side of equation.

4Fe + 3O₂ —> 2Fe₂O₃

4 Fe on each side

6 O on each side

Topic 3: Mathematics of Chemistry

Gram Formula Mass (GFM), Formula Mass, Molecular Weight, Molecular Mass - the sum of the atomic masses of each element within the compound

Percent Composition (On Table T) - (mass of part / mass of whole compound) * 100

Mole Calculations (On Table T)

Moles = Mass / GFM

Mass = (Moles)*(GFM)

Stoichiometry / Ratios

The coefficients in a balanced equation represent moles. Find the ratio between two compounds: this is a molar ratio.

CH₄ + 2O₂ —> CO₂ + 2H₂O

CH₄ and H₂O are in a 1:2 ratio

3 moles of CH₄ will produce 6 moles of H₂O

Molecular Formulas from Empirical Formulas

(Gram Formula Mass) / (Weight of Empirical Formula)

Multiple this ratio by the empirical formula to get the molecular formula

Example:

Empirical Formula: CH₂ and GFM = 28 g/mol

C 1×12 = 12

H 2×1 = 2

12 + 2 + 14 g/mol

(28) / (14) = 2

2(CH₂) = C₂H₄

Topic 4: Physical Behavior of Matter

Phases of Matter

Solid: definite shape and volume, rigid, fixed patterns

Liquids: indefinite, definite volume

Gases: indefinite shape and volume, fills container completely

Heating and Cooling Curves

Lines that are increasing (or decreasing) in temperature

Kinetic Energy is changing

Lines that are flat

Kinetic Energy is remaining the same, but Potential Energy is changing

Freezing: liquid to a solid, requires heat of fusion (334 J)

Melting: liquid to a solid, requires heat of fusion (334 J)

Condensation: gas to a liquid, requires heat of vaporization (2260 J)

Boiling/Vaporization: liquid to a gas, requires heat of vaporization (2260 J)

Sublimation: solid to a gas

Deposition: gas to a solid

Temperature is AVERAGE KINETIC ENERGY

Temperature is not affected by amount, highest temperature = highest average kinetic energy

Heat of Fusion: amount of heat needed to melt/freeze

Heat of Vaporization: amount of heat needed to vaporize/condense

Gas Law Relationships

PV / T = PV / T

(P-T-V Chart)

If something is constant or not given, remove it from the equation

Graphs:

Pressure and Volume = Inverse

Pressure and Temperature = Direct

Volume and Temperature = Direct

Ideal vs Real Gases

Volume is negligible in ideal gases

Ideal gases have no attraction

Ideal gases may not transfer energy between each other

Particles move in random, straight-line motion

Gases are more ideal at Higher Temperature and Lower Pressure (HiTLoP - high temp, low pressure)

Avogadro’s Law

Two gases that occupy the same volume have the same temperature, same pressure, and same number of molecules

The masses of the gases may be different based on their identity and GFM

Topic 5: The Periodic Table

Mendeleev - organized elements by atomic mass

Modern Periodic Law - elements are arranged by atomic number

Organization

Groups - Columns (18)

Periods - Rows (7)

Group 1 = Alkali Metals

Group 2 = Alkaline Earth Metals

Groups 3-12 = Transition Metals

Group 17 = Halogens

Group 18 = Nobel Gases

Metals (left side of table, most elements)

Non-Metals (right side of table, and hydrogen)

Metalloids (touching the diagonal except Al and Po)

Liquids = Hg, Br

Gases = H₂, O₂, N₂ F_2, Cl₂, He, Ne, Ar, Rn, Xe

The rest are solid

Properties

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

Different Molecular Structures = Different Properties

Trends

Atomic Radius: size of the atom

Across a period: radius DECREASES

Down a group: radius INCREASES

Ionization Energy: amount of energy needed to remove electrons in an atom

Across a period: IE INCREASES

Down a group: IE DECREASES

Electronegativity: an atoms attraction for electrons

Across a period: EN INCREASES

Down a group: EN DECREASES

Groups

Groups are organized by the number of valence electrons each atom possesses

Group 1 has 1 valence electron, Group 13 has 3, Group 15 has 5, etc..

Elements within the same group have similar chemical properties due to the same number of valence electrons.

Periods

Organized by number of electron shells, periods do NOT have similar properties

Topic 6: Bond Polarity

Bond Energy

Breaking Bonds —> Absorbs Energy

Forming Bonds —> Releases Energy

B.A.R.F

Break Absorb, Release Form

Types of Bonds

Covalent: between 2 or more non-metals (sharing of electrons)

Ionic: between a metal and a non-metal (transfer of electrons)

Metallic: just a metal (sea of mobile electrons)

Types of Solids/Properties

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

Lewis Dot Structures

Show only valence electrons of atom, the most dots an atom can have is 8 (Noble Gas Configuration)

Covalent Structures: between 2 non-metals atoms

Each line = 2 electrons (a pair of electrons)

Ionic Structures: between metal and non-metal

Have brackets, not lines

Both ions must have charges present

Metals have zero dots and a positive charge

Non-Metals have eight dots in a bracket and a negative charge

Ions form to achieve Noble Gas Configuration

All diatomic elements are Non-Polar with Non-Polar bonds

H₂O and NH₃ are Polar because of lone pairs

CCl₄, CO₂, and CH₄ are Non-Polar with Polar Bonds

Bond Polarity

S.N.A.P

Symmetrical = Non-Polar, Asymmetrical = Polar

Polar: unequal distribution of charge, one side has electrons more often (the more electronegative side)

Polar Bond: 2 different elements

Non-Polar: equal distribution of charge, electrons are shared equally among all sides of the molecule

Non-Polar Bond: same element on both sides of the bond

Intermolecular Forces

Between covalent compounds ONLY

Non-Polar molecule = London Dispersion Forces (weakest)

Polar Molecules = Dipole-Dipole Forces

H₂O, NH₃, HF (H- with F, O, or N) = Hydrogen Bonding (strongest)

Intermolecular forces determine melting point/boiling point

Weaker forces = lower melting/boiling point

Strongest forces = higher melting/boiling point

Topic 7: Solutions

Formulas (Table T)

Moles = mass / GFM

Molarity = moles / liters

Parts Per Million = (part/whole) * 1,000,000 or 1×10⁶

Solute: substance that is being dissolved

Solvent: substance that is dissolved into (usually water)

Solution: solvent + solute

Solubility Factors

Temperature: increase temperature = increase solubility (NOT gases)

Gases: Increase temperature = decrease solubility

Pressure (only gases): increase pressure = increase solubility

Nature of Solvents & Solutes: “like dissolves like”

Polar (water) dissolves Polar

Non-Polar dissolves Non-Polar

Ionic substances dissolve in Polar solvents

Table G (Solubility Curve)

Read in 100g of water, alter as necessary

If 200g of water = double amount from table

If 50g of water = half the amount from the table

A solubility reading ABOVE the line = SUPERSATURATED

A solubility reading ON the line = SATURATED

A solubility reading BELOW the line = UNSATURATED

When given 2 values to read, always subtract the values

If solute remain on bottom of beaker: Saturated solution (at equilibrium)

If solute dissolved: Unsaturated solution

If solute causes precipitation of more particles: Supersaturated Solution

Colligative Properties

Adding a solute to a solvent LOWERS the freezing point and RAISES the boiling point

The more ions it breaks into, or the higher concentration, the more effect it has on the freezing point and boiling point

Covalent compounds do NOT break up

Vapor Pressure (Table H)

Shows vapor pressure of substances at given temperatures

Increase temperature = Increase vapor pressure (for any substance, direct relationship)

Boiling: when atmospheric pressure = vapor pressure (can be any temperature)

Lower VP: Stronger Intermolecular Forces and Higher Boiling Point (Ethanoic Acid)

Higher VP: Weaker Intermolecular Forces and Lower Boiling Point (Propanone)

Topic 8: Kinetics and Equilibrium

Collision Theory

Collisions between particles must have proper energy and proper orientation

which make effective collisions occur

Factors that Affect Rate of a Reaction

Concentration: higher concentration = faster rate and more effective collisions

Temperature: higher temperature = more collisions and faster rate

Surface Area: more surface area = more area for more effect collisions

Pressure (only gases): higher pressure = more collisions

Nature of Reactants: covalent reacts slower; ionic reacts faster

Catalyst: lowers activation energy by providing an alternate pathway

Potential Energy Diagrams

Anything that starts from the bottom (zero) are potential energies

Reactants: left side

Products: right side

Forward Activation Energy: left side

Reverse Activation Energy: right side

Heat of Reaction (ΔH) = PE Products - PE Reactants

Equilibrium

Rates are Equal and Concentration is Constant

Phase Equilibrium: same compound or element but in different phases

Solution Equilibrium: same compound or element but in one phase is in solution

Chemical Equilibrium: a state in which the rate of the forward reaction equals the rate of the backward reaction

On a graph: flat lines mean concentration is remaining constant

LeChatlier’s Principle

Increase Concentration or Temperature: shift to opposite side of what you’re increasing

Decrease Concentration or Temperature: shift to same side of what you’re decreasing

Increase Pressure: shift to side with less moles of gas

Decrease Pressure: shift to side with more moles of gas

Side shifting toward: concentration increases

Side shifting away from: concentration decreases

Forward Reaction: left to right

Reverse Reaction: right to left

Catalysts: speed up reactions, but do not increase concentration

Entropy (disorder)

Solids have the least entropy, gases have the most entropy

The side with more moles has entropy

Increase Temperature = Increase Entropy

Nature prefers lower energy and higher entropy

Topic 9: Redox and Electrochemistry

Oxidation

Electrons appear on product (right) side

Oxidation number INCREASES

Electrons DECREASE

Reduction

Electrons appear on reactant (left) side

Oxidation number DECREASES

Electrons INCREASE

Oxidation Numbers

(follow rules in order of priority)

All uncombined elements (not in a compound) have an oxidation number of zero

Ions in a compound have an oxidation numbers equal to their charges

Group 1 metals = +1, Group 2 metals = +2, Fluorine = -1, (Cl, Br, I are too unless with something more electronegative)

H = +1 (unless ONLY with a metal then its -1)

O = -2 (with Fluorine O = +2, in peroxide O₂^-2 = -1)

Sum of all oxidation numbers in compounds = 0 or stated charge if a polyatomic ion

Redox Reactions

Single Replacement, Synthesis and Decomposition are Redox Reactions

Any reaction where an element changes oxidation numbers

Look for single elements somewhere

Double Replacements are NEVER redox

Table J (Activity Series)

Single Element must be higher than element in compound to replace (occur spontaneously)

Higher elements are oxidation (anode), lower elements are reduction (anode)

Mnemonics

O.I.L.R.I.G

OIL - Oxidation is Loss (of electrons)

RIG - Reduction is Gain (of electrons)

An Ox - Anode Oxidation

Red Cat - Reduction Cathode

VAN - Voltaic cell Anode is Negative

APE - Anode is Positive in Electrolytic cell

Voltaic (Electrochemical) Cell

Spontaneous, require no external energy

Chemical Energy —> Electric Energy

Salt Bridge: migration of ions

Wire: migration of electrons

Electrons flow from Anode to Cathode

(Away from Anode)

Anode = negative (-)

Half Reaction

Show only the Oxidation or Reduction

MUST show charges

Combined together (with balanced and canceled electrons) = full redox

The element not changing oxidation number from side to side = spectator (cross it out)

Electrolytic Cell

Non-Spontaneous, requires a battery

Electrical Energy —> Chemical Energy

NO Salt Bridge

Electrons flow from Anode to Cathode

(Away from Anode)

Used for Electroplating

Positive Ions flow to negative cathode

Negative Ions flow to positive anode

Agents

Oxidizing Agent: causes oxidation, so it’s reduced

Reducing Agent: causes reduction, so it oxidized

Topic 10: Acids, Bases and Salts

Table K - Common Acids

Table L - Common Bases

Table M - Indicators

Acids: Begin with H

Bases: Metal + OH

Properties

Acids: taste sour, turns litmus red, pH less than 7

Bases: feel slippery, taste bitter, turns litmus blue, pH greater than 7

Both: conduct electricity, react with each other to form salt and water

Electrolytes: acids, bases and ionic salts (substances that conduct electricity)

Acid/Base Theory

Arrhenius Theory

Acid: substance that donate H⁺ as it’s only positive ion

Base: substance that donates OH^- as it’s only negative ion

Bronsted-Lowry Theory (Alternate Theory)

Acid: substance that donates an H⁺ ion

Base: substance that accept an H⁺ ion

(form conjugate pairs)

Acid has 1 more H⁺

Base has 1 less H⁺

Increase [H⁺], Decrease [OH^-] = more acidic, lower pH

Decrease [H⁺], Increase [OH^-] = more basic, higher pH

Neutralization Reaction

acid + base —> salt + water

(double replacement reaction)

Titration

Adding measured volumes of a known acid/base to an unknown concentration of acid/base to reach neutralization

M_AV_A = M_BV_B

(M = molarity, V = volume)

Don’t forget to count the H’s and OH’s

pH

Scale of 1-14

pH 1-6 = Acid

pH 7 = Neutral

pH 8-14 = Base

Each power of 10 for concentration = 1 pH unit

ex:

pH = 7 hundredfold increase hydronium —> pH 5 (10²)

Indicators

Change color through specific pH ranges (Table M)

Topic 11: Organic Chemistry

All organic compounds contain carbon because of it’s ability to bond to itself multiple times

Saturated Compounds - all single C-C bonds (alkanes)

Unsaturated Compounds - at least 1 or double C-C bonds (alkenes and alkynes)

Naming Compounds

Name longest carbon chains with Table P and Table Q

Use number to identify which carbona group is hanging off o

A carbon group can NOT be put on the first or last carbons; it will just extend the longest chain

Functional Groups (Table R)

Look for “O” or “N” and what’s around it

Distinguish between CO, COO, COOH and OH

Circle C_xH_y groups

Isomers

Same molecular formula (same GFM), different structural formula

Different Structures = Different Properties = Different Functional Groups

Organic Reactions

Addition: with UNSATURATED hydrocarbons (alkene and alkyne)

Substitution: with SATURATED hydrocarbons (alkane)

Combustion: always produces CO₂ AND H₂O

Esterification: Alcohol + Organic Acid —> Ester + Water

Fermentation: Sugar (C₆H₁₂O₆) —> Alcohol + Carbon Dioxide

Saponification: Ester + Base —> Acid + Alcohol (process of making soap)

Polymerization: Repeating units —> Polymers (involves smaller molecules joining to make one big molecule)

Cracking: long chain hydrocarbon —> small chain hydrocarbon

Topic 12: Nuclear Chemistry

Table N - Half Lives

Table O - Decay Symbols

Particles

Alpha: heaviest particle (greatest mass), positive charge, attracted to negative plate (low energy), least penetrating power

Beta: no mass, negative charge, attracted to positive plate

Gamma: no mass, no charge, neutral (high energy), greatest penetrating power

Transmutation

When one atomic nuclei is changed into the nucleus of a different element

Natural Transmutation (one reactant)

Artificial Transmutation (two reactants)

Fission (releases energy)

One atom absorbs a neutron and splits into two or more pieces, giving off a tremendous amount of energy

Fusion (releases more energy)

When the two light nuclei unite to form a heavier nucleus, fusion creates more energy than fission

Balancing Nuclear Reactions

Set masses of both sides equal

Set Atomic Numbers of both sides equal

Half Life

Time it takes for half of a sample to decay

Divide amount in half for each half life

Uses of Radioisotopes

C-14: used in radioactive dating of once living organisms

U-238: used in radioactive dating of geological (non-living substances)

I-131: used to treat thyroid disorders

Co-60: used in radiation therapy for cancer treatment