Chemistry : Ultimate Review Guide

Unit 2 : Types & Properties of Matter

Types of Matter — 4-8% (4-8% of the PreComp will be on this)

Matter—something that has mass and volume

• mass: amount of matter in an object

• volume: amount of space an object takes up

Elements—one type of atom

• can exist as atoms or molecules

Molecules—two or more atoms chemically bound together

• same or different types of atoms

Compounds—two or more different elements

• have a set ratio of elements

Properties of Matter — 4-8%

Elements can’t be broken down using chemistry

• Identity determined by # of protons in nucleus

• Isotopes determined by # of neutrons in nucleus

Compounds can be broken down using chemistry

• Has a fixed chemical composition throughout

• Made up of two or more different elements chemically combined

Mixtures contain two or more substances

• Homogeneous Mixture—one or more substances dissolved in another substance

  • Solutions

    • Solute—the substance being dissolved

      • Major Component

    • Solvent—the substance doing the dissolving

      • Major Component

    • aqueous (aq) = “dissolved in water”

• Heterogeneous Mixture—mixture of substances that remain physically separate

  • Suspensions—contains large particles that settle out of a mixture

    • If you can see individual particles, then it’s a suspension

Separation Techniques — 0-4%

Filtration separates solids from liquids and gases

• A filter only allows fluid to pass through, leaving solids behind

• Filtrate—the fluid that passes through the filter

  • Cannot be used on solutions (heterogeneous mixtures only)

Distillation separates liquids based on their different boiling points

• A mixture of fluids is boiled

  • Fluid with lowest BP evaporates first—it has the weakest IMFs

• Vapor of lower BP fluid cools and condenses into another container

Chromatography separates liquids based on their solubilities

• A drop of the mixture goes on a stationary phase

  • Stationary Phase—stays in place

• The mobile phase travels over the stationary phase

  • Mobile Phase—solvent

States of Matter — 0-4%

Solids have definite shape and volume

• Low energy—vibrate in place

• Regular particle pattern—touching

Liquids have definite volume and take the shape of their container

• Some energy—vibrate and slide past each other

• Irregular particle pattern—touching

Gases take the shape and volume of their container

• High energy—vibrate, move quickly, bounce off of each other

• Irregular particle pattern—not touching, as spread out as possible

• Compressible

Changes in Matter

• Solid→Liquid—Melting

• Liquid→Solid—Freezing

• Gas→Liquid—Condensation

• Liquid→Gas—Boiling

• Solid→Gas—Sublimation

• Gas→Solid—Deposition

Temperature, Heat, & Heating Curve — 4-8%

Heat transfers from one substance to another

• Exothermic—system releases heat to surroundings

• Endothermic—system absorbs heat from surroundings

Temperature—measure of the average kinetic energy of particles in a substance

Heat—energy transferred from one system to another as a result of a difference in temperature (only exists in in the process)

Temperature Conversions

• Fahrenheit to Celsius

  • Cº = 5/9(Fº-32)

• Celsius to Fahrenheit

  • Fº = 9/5(Cº+32)

• Celsius to Kelvin

  • Cº = K-273

• Kelvin to Celsius

  • K = Cº+273

Modes of Transfer

• Convection—energy transfer due to the bulk motion of fluids of different temps

• Conduction—energy transfer due to the difference in temperature in adjoining regions (transfer through particle collisions)

• Radiation—transfer of energy through electromagnetic waves

Heating Curves

Unit 3 : Periodic Table & Trends

Classification & Families of Elements — 4-8%

Periods go across the periodic table (left & right)

• Elements in the same period have the same number of electron shells

  • # of Shells = Period #

Groups go down the periodic table (up & down)

• Elements in the same group have similar properties and same number of valence electrons

  • # of V.E^- = Group #

Characteristics of Metals

• Good conductors, lustrous (shiny), malleable, ductile, high melting points, form cations

• Alkali Metals

  • soft, highly reactive, form +1 ions

• Alkaline Earth Metals

  • soft, very reactive, form +2 ions

• Transition Metals

  • most commonly known metals, often have very colorful ions, form ions with a variety of charges

• Halogens

  • diatomic elements, often gaseous at room temp

• Noble Gases

  • inert/unreactive gases, characteristically light up when attached to electricity

• Lanthanides & Actinides

  • reactive with halogens, actinides are radioactive, rare earth metals

Characteristics of Nonmetals

• Insulators, dull, brittle, low melting points, form anions

Subatomic Particles & Ions — 4-8%

Isotopes are atoms of the same element with a different number of neutrons

Ions are formed when atoms give up or gain electrons

• When electrons are gained, a negative ion is formed

• When electrons are lost, a positive ion is formed

Radius, Ionization Energy, Electronegativity — 8-16%

Atomic Radius—the distance from an atom’s nucleus to its outermost electrons

• Goes from top right to bottom left

• Cs is actually larger than Fr—Cesium is the largest element, not Francium

Ionization Energy—the energy required to remove an electron from a neutral atom in its gaseous state

Electronegativity—an atom’s ability to attract shared electrons in a chemical bond

• Polar Bonds—elements have a high difference in electronegativity

• Nonpolar Bonds—elements have a low difference in electronegativity

  • C-H bonds are nonpolar

  • Same element bonds are nonpolar

Unit 4 : Chemical Bonding

Ionic Bonds — 8-20%

Metal & Nonmetal—electrons are transferred

• electrostatic attraction between oppositely charged particles

cation (+), metal or polyatomic ion

anion (-), nonmetal or polyatomic ion

Properties of Ionic Bonds

• high melting points

• solid does not conduct electricity

• both liquid & solution will conduct electricity

• some are soluble in water

Ionic bonds must follow the rule of zero charge

Covalent Bonds — 12-28%

Two Nonmetals—electrons are shared

• each atom contributes 1 bond

Properties of Covalent Bonds

• low melting points

• do not conduct electricity

• some dissolve in water

Metallic Bonds — 4-8%

Two Metals—electrons are delocalized

• atoms are surrounded by a “sea” of shared electrons

Properties of Metallic Compounds

• high melting points

• do not dissolve in water

• conduct electricity as both a liquid & solid

Unit 5 : Molar Mass

Molar Mass — 4-8%

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

1 Mole = 6.022 × 10²³ particles (Avogadro’s number)

Molar mass is the sum of the atomic masses of all the atoms in a formula.

Calculating Molar Mass

• For a molecule, add up the molar masses of all elements in the compound.

• Example: Molar mass of H₂O = 2(1.008 g/mol) + 16.00 g/mol = 18.016 g/mol.

Moles & Conversion

1 Mole of a substance = mass (g) ÷ molar mass (g/mol).

Use stoichiometric relationships to convert between moles, mass, and volume (for gases).

Percent Composition — 4-4%

Percent composition is the mass percent of each element in a compound.

Formula

Percent Composition = (Mass of element ÷ Mass of compound) × 100

Example:

• For NaCl, the percent composition of Na is:

(22.99 g/mol ÷ 58.44 g/mol) × 100 ≈ 39.3%.

• The percent composition of Cl is:

(35.45 g/mol ÷ 58.44 g/mol) × 100 ≈ 60.7%.

Empirical Formula — 4-12%

The empirical formula represents the simplest whole-number ratio of elements in a compound.

How to Find the Empirical Formula

1. Convert the mass of each element to moles.

2. Divide each element’s mole value by the smallest number of moles.

3. Round to the nearest whole number if needed.

Example:

• For a compound with 40.0 g C and 6.7 g H:

1. Convert to moles:

• C: 40.0 g ÷ 12.01 g/mol = 3.33 mol

• H: 6.7 g ÷ 1.008 g/mol = 6.64 mol

2. Divide by the smallest mole number (3.33):

• C: 3.33 ÷ 3.33 = 1

• H: 6.64 ÷ 3.33 ≈ 2

Empirical formula: CH₂

Molecular Formula — 4-8%

The molecular formula is the actual number of atoms of each element in a compound. It may be the same as the empirical formula or a multiple of it.

How to Find the Molecular Formula

1. Calculate the empirical formula mass (EFM).

2. Divide the molar mass of the compound by the EFM.

3. Multiply the empirical formula by this factor.

Example:

• If the empirical formula is CH₂ and the molar mass of the compound is 56.08 g/mol,

EFM = 12.01 + 2(1.008) = 14.026 g/mol.

56.08 ÷ 14.026 ≈ 4.

Thus, the molecular formula is C₄H₈.

Unit 6 : Chemical Reactions & Stoichiometry

Acid-Base Reactions

• An acid (H^- ion) and a base (OH⁺ ion) react to form a salt and a water.

Precipitation Reactions & Solubility — 4-8%

Precipitation Reaction:

A reaction where two aqueous solutions mix and an insoluble solid (called a precipitate) forms and settles out.

Solubility

• Soluble: Substances that dissolve well in water (form aqueous solutions).

• Insoluble: Substances that do not dissolve well and form solids (precipitates).

How to Predict a Precipitate

1. Write the formulas of the reactants and possible products.

2. Use solubility rules to check if any product is insoluble.

3. If an insoluble product forms, that’s the precipitate.

Common Solubility Rules

• Nitrates (NO₃⁻) and acetates (CH₃COO⁻) are always soluble.

• Alkali metals (Group 1) compounds are soluble.

• Halides (Cl⁻, Br⁻, I⁻) are soluble except with Ag⁺, Pb²⁺, Hg₂²⁺.

• Sulfates (SO₄²⁻) are soluble except with Ba²⁺, Pb²⁺, Ca²⁺, Sr²⁺.

• Most carbonates (CO₃²⁻), phosphates (PO₄³⁻), hydroxides (OH⁻) are insoluble except with alkali metals and NH₄⁺.

Redox Reactions — 4-8%

Redox (Oxidation-Reduction) Reactions:

Chemical reactions where electrons are transferred between substances.

Key Terms

• Oxidation: Loss of electrons (increase in oxidation state)

• Reduction: Gain of electrons (decrease in oxidation state)

Oxidation Numbers (Oxidation States)

• An oxidation number is a number assigned to an element in a compound that shows how many electrons it has gained, lost, or shared compared to its neutral atom.

• It helps keep track of electron transfer in redox reactions.

Basic Rules for Assigning Oxidation Numbers

1. The oxidation number of any pure element (like O₂, N₂, or Fe metal) is 0.

2. For a simple ion, the oxidation number equals the charge of the ion (e.g., Na⁺ is +1, Cl⁻ is -1).

3. Oxygen usually has an oxidation number of -2 (except in peroxides where it’s -1).

4. Hydrogen usually has an oxidation number of +1 when bonded to nonmetals, and -1 when bonded to metals.

5. The sum of oxidation numbers in a neutral compound is 0.

6. The sum of oxidation numbers in a polyatomic ion equals the ion’s charge.

Combustion Reactions — 4-8%

Combustion Reaction:

A chemical reaction where a substance (usually a hydrocarbon) reacts rapidly with oxygen (O₂) to produce carbon dioxide (CO₂) and water (H₂O), releasing heat and light.

Decomposition Reactions

Rules for Decomposition Reactions

1. Binary Compounds often break down into their elements:

   AB = A + B

   • 2HgO = 2Hg + O₂

2. Metal Carbonates decompose into metal oxides and carbon dioxide:

   MCO₃ = MO + CO₂

   CaCO₃ = CaO + CO₂

3. Metal Hydroxides decompose into metal oxides and water:

   M(OH)_n = MO + nH₂O

   2NaOH = Na₂O + H₂O

4. Metal Chlorates decompose into metal chlorides and oxygen gas:

   MClO₃ = MCl + O₂

Limiting Reactant & Yield

Actual Yield = (Percent Yield/100) x Theoretical Yield (how much is actually produced)

Percent Yield = Actual/Theoretical x 100

Theoretical Yield = grams of limiting reactant → grams of prdouct

Limiting Reactant: The reactant that is completely used up in a chemical reaction and limits the amount of product formed.

• to find: convert moles of each reactant to moles of product

  • whichever one is less is the limiting!

Excess Reactant: The reactant that is not completely used up and remains after the reaction is complete.

Steps to Find the Limiting Reactant

1. Convert the given mass (or moles) of each reactant to moles of product using stoichiometry.

2. Compare the moles of product each reactant can produce.

3. The reactant that produces less product is the limiting reactant.

4. The other is the excess reactant.

Gas Laws

Gas Laws

Boyle’s Law:

• P₁V₁ = P₂V₂

Charles’s Law:

• V_1/T₁ = V_2/T₂

Gay-Lussac’s Law:

• P_1/T₁ = P_2/T₂

Avogadro’s Law:

• V_1/n₁ = V_2/n₂

Combined Gas Law:

• P₁V₁/T₁ = P₂V₂/T₂

Ideal Gas Law:

• PV = nRT

      n = number of moles of gas

      R = the ideal gas constant, usually 0.0821 L·atm/mol·K

      T = temperature in Kelvin