Chapter 1 — Matter, Measurement & Problem Solving

What is Matter?

Matter is defined as anything that has mass and occupies space. All physical objects, from air to water to a pencil, are considered matter. Matter is composed of tiny particles called atoms.

• Examples: Water, air, your phone, a pencil, an atom

• Atoms: The smallest basic unit of an element

Atoms vs. Molecules

An atom is the smallest unit of an element, while a molecule consists of two or more atoms chemically bonded together. The type and arrangement of atoms in a molecule determine its physical and chemical properties.

• Atom Example: H (hydrogen atom), O (oxygen atom)

• Molecule Example: H2O (water), O2 (oxygen gas)

Scientific Approach

• Hypothesis: A testable explanation or prediction (e.g., "If I increase the temperature of water, then it will evaporate faster.")

• Scientific Law: Describes a consistent pattern in nature (what happens)

• Theory: A well-supported explanation for why something happens (not a guess)

Classifying Matter

• Pure Substance: Fixed composition; can be an element (one type of atom, e.g., gold, oxygen) or a compound (two or more elements chemically bonded, e.g., H2O, CO2)

• Mixture: Two or more substances physically combined

  • Homogeneous mixture: Uniform composition (e.g., salt water, air)

  • Heterogeneous mixture: Non-uniform composition (e.g., salad, oil and water)

States of Matter

• Solid: Fixed shape and volume; particles vibrate in place

• Liquid: Fixed volume, no fixed shape; particles move around each other

• Gas: No fixed shape or volume; particles are far apart and move freely

Crystalline vs. Amorphous Solids

• Crystalline: Organized, repeating pattern (e.g., salt, quartz)

• Amorphous: No long-range order (e.g., glass, some plastics)

Separating Mixtures

• Filtration: Separates solids from liquids using a filter

• Distillation: Separates substances based on boiling points

• Decanting: Carefully pouring off a liquid from a solid

Physical vs. Chemical Properties

• Physical Property: Observed without changing identity (e.g., color, density, melting point)

• Chemical Property: Describes how a substance can change into another (e.g., flammability, reactivity)

Physical vs. Chemical Changes

• Physical Change: Substance remains the same (e.g., melting ice, boiling water)

• Chemical Change: New substance is formed (e.g., burning wood, rusting iron)

Energy in Chemistry

• Kinetic Energy: Energy of motion

• Potential Energy: Stored energy due to position or arrangement

• Thermal Energy: Energy from particle motion; related to temperature

SI Units

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: Kelvin (K)

• Amount: mole (mol)

• Electric current: ampere (A)

• Luminous intensity: candela (cd)

Metric Prefixes

Density

Density is the amount of mass per unit volume. The formula is:

• D = density

• m = mass

• V = volume

Example: If a substance has a mass of 20 g and a volume of 5 mL,

Intensive vs. Extensive Properties

• Intensive: Does not depend on amount (e.g., density, temperature)

• Extensive: Depends on amount (e.g., mass, volume)

Accuracy vs. Precision

• Accuracy: How close a measurement is to the true value

• Precision: How close repeated measurements are to each other

Significant Figures

• Nonzero digits are always significant

• Zeros between nonzero digits are significant

• Leading zeros are not significant

• Trailing zeros after a decimal are significant

Multiplication/Division: Use the fewest significant figures

Addition/Subtraction: Use the fewest decimal places

Chapter 2 — Atoms & Elements

Brownian Motion

Brownian motion is the random movement of particles suspended in a fluid, observed by Robert Brown and explained by Einstein as resulting from collisions with moving molecules. This provided evidence for the particulate nature of matter.

Early Atomic Theory

• Leucippus & Democritus: Proposed matter is made of indivisible atoms

• Aristotle & Plato: Believed matter could be divided infinitely

Three Important Laws

• Law of Conservation of Mass: Matter is not created or destroyed in chemical reactions

• Law of Definite Proportions: Compounds have elements in fixed ratios

• Law of Multiple Proportions: Elements combine in small whole-number ratios to form different compounds

Dalton's Atomic Theory

• Matter is made of atoms

• Atoms of the same element are similar

• Atoms of different elements are different

• Compounds form from whole-number ratios of atoms

• Chemical reactions rearrange atoms

Subatomic Particles

Protons determine the element's identity.

Atomic Number and Mass Number

• Atomic Number: Number of protons

• Mass Number: Protons + Neutrons

• Neutrons: Mass Number - Protons

Example: Carbon-14: 6 protons, 8 neutrons, 6 electrons (if neutral)

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons (and thus different mass numbers).

• Example: Carbon-12 (6 neutrons), Carbon-14 (8 neutrons)

Ions

• Cation: Positive ion (loses electrons, e.g., Na → Na+)

• Anion: Negative ion (gains electrons, e.g., Cl + e- → Cl-)

Periodic Table Groups

• Group 1 (Alkali Metals): 1 valence electron, form +1 ions

• Group 2 (Alkaline Earth Metals): 2 valence electrons, form +2 ions

• Group 17 (Halogens): 7 valence electrons, form -1 ions

• Group 18 (Noble Gases): Full outer shell, unreactive

Main-group properties are strongly related to valence electrons.

Metals vs. Nonmetals

• Metals: Left side of the periodic table, tend to lose electrons

• Nonmetals: Right side, tend to gain or share electrons

• Transition Metals: Middle of the periodic table

Atomic Mass vs. Molar Mass

• Atomic Mass: Average mass of an element's isotopes (amu)

• Molar Mass: Mass of one mole of a substance (g/mol)

Example: Carbon: atomic mass ≈ 12.01 amu, molar mass ≈ 12.01 g/mol

The Mole and Avogadro's Number

• 1 mole = particles (Avogadro's number)

• Used as a counting unit in chemistry

Mole Conversions

• Grams to Moles:

• Moles to Grams:

• Moles to Particles:

• Particles to Moles:

Chapter 3 — Molecules & Compounds

Ionic vs. Covalent Bonds

• Ionic Bond: Metal + Nonmetal; electrons are transferred (e.g., NaCl)

• Covalent Bond: Nonmetal + Nonmetal; electrons are shared (e.g., H2O)

Ionic compounds form lattices of alternating cations and anions.

Ionic vs. Molecular Compounds

• Ionic Compound: Contains ions (e.g., NaCl); smallest unit is a formula unit

• Molecular Compound: Atoms connected by covalent bonds (e.g., CO2); smallest unit is a molecule

Molecular Elements

Some elements exist naturally as molecules, especially the seven diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2.

• Memory trick: Have No Fear Of Ice Cold Beer

Empirical vs. Molecular Formula

• Empirical Formula: Simplest whole-number ratio (e.g., CH2O)

• Molecular Formula: Actual number of atoms in a molecule (e.g., C6H12O6)

Structural Formula

Shows how atoms are connected (e.g., H–O–H for water).

Naming Ionic Compounds

• Metal + Nonmetal: Name metal first, change nonmetal ending to -ide (e.g., NaCl = sodium chloride)

• Transition metals: Use Roman numerals for charge (e.g., FeCl2 = iron(II) chloride)

Naming Molecular Compounds

• Use prefixes to indicate number of atoms (e.g., CO2 = carbon dioxide, N2O5 = dinitrogen pentoxide)

Formula Mass

To find formula mass:

1. Find atomic mass of each element

2. Multiply by number of atoms

3. Add together

Example: H2O: (2 × 1.01) + 16.00 = 18.02 g/mol

Mass Percent

Mass percent is the percentage of a compound's mass from one element:

Example: If oxygen is 16 g in a 20 g compound:

Finding an Empirical Formula

1. Convert grams to moles

2. Divide by the smallest number of moles

3. Get whole numbers

4. Write the empirical formula

Example: C = 1 mol, H = 2 mol, O = 1 mol → CH2O

Combustion Analysis

Combustion analysis determines the formula of a compound by burning it and measuring the amounts of CO2 and H2O produced. This allows calculation of the empirical formula using mole ratios.

Key Concepts to Memorize

• Matter vs. substance

• Element vs. compound

• Homogeneous vs. heterogeneous

• Solid/liquid/gas

• Physical vs. chemical properties/changes

• Density, SI units, metric prefixes

• Significant figures, accuracy vs. precision

• Dimensional analysis

• Proton/neutron/electron, atomic number, mass number, isotopes, ions

• Periodic table groups, valence electrons, metals vs. nonmetals

• Atomic mass, molar mass, Avogadro's number, mole conversions

• Ionic vs. covalent, naming compounds, empirical/molecular formulas, formula mass, mass percent, combustion analysis