Chemistry BASIS 8 : Comp

_{wow last one!! goodluck on precomps}

_{Unit 1 will not be featured on the PreComp or Comp}

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

Write the formulas of the reactants and possible products.
Use solubility rules to check if any product is insoluble.
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

The oxidation number of any pure element (like O₂, N₂, or Fe metal) is 0.
For a simple ion, the oxidation number equals the charge of the ion (e.g., Na⁺ is +1, Cl⁻ is -1).
Oxygen usually has an oxidation number of -2 (except in peroxides where it’s -1).
Hydrogen usually has an oxidation number of +1 when bonded to nonmetals, and -1 when bonded to metals.
The sum of oxidation numbers in a neutral compound is 0.
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

Binary Compounds often break down into their elements:
AB = A + B
- 2HgO = 2Hg + O₂
Metal Carbonates decompose into metal oxides and carbon dioxide:
MCO₃ = MO + CO₂
CaCO₃ = CaO + CO₂
Metal Hydroxides decompose into metal oxides and water:
M(OH)_n = MO + nH₂O
2NaOH = Na₂O + H₂O
Metal Chlorates decompose into metal chlorides and oxygen gas:
MClO₃ = MCl + O₂

Limiting Reactant & Yield

Actual Yield = (Percent Yield/100) x Theoretical Yield

Percent Yield = Actual/Theoretical x 100

Theoretical Yield = Limiting Reactant (In Moles) x Moles of Product/Moles of Reactant x Molar Mass of Product

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

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

Steps to Find the Limiting Reactant

Convert the given mass (or moles) of each reactant to moles of product using stoichiometry.
Compare the moles of product each reactant can produce.
The reactant that produces less product is the limiting reactant.
The other is the excess reactant.

Gas Laws

Boyle’s Law:

P₁V₁ = P₂V₂

Charles’s Law:

V₁T₁ = V₂T₂

Gay-Lussac’s Law:

P₁T₁ = P₂T₂

Avogadro’s Law:

V₁n₁ = V₂n₂

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