Chapter 1: Chemistry Basics - Matter and Measurement

Matter

Matter is anything that has mass and occupies space. It is the fundamental substance of chemistry.

• Definition: Matter includes all solids, liquids, and gases.

• Example: Water, air, and iron are all forms of matter.

Classification of Matter

Matter can be classified using a flow chart into pure substances and mixtures.

• Pure Substance: Has a fixed composition; can be an element or a compound.

• Mixture: Contains two or more substances physically combined; can be homogeneous or heterogeneous.

• Homogeneous Mixture: Uniform composition throughout (e.g., salt water).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

Elements and Atoms

An element is a pure substance made of only one kind of atom. An atom is the smallest unit of an element that retains its properties.

• Example: Oxygen (O) is an element; a single oxygen atom is its smallest unit.

Periodic Table of Elements

The periodic table organizes elements by increasing atomic number and groups them by similar properties.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows; elements in the same period have the same number of electron shells.

• Chemical Symbol: One or two letters representing an element (e.g., Na for sodium).

• Metals: Groups 1A and 2A; transition metals used in certain problems.

• Halogens: Group 7A (17); highly reactive nonmetals.

• Noble Gases: Group 8A (18); inert gases.

• Nonmetals: Located on the right side of the periodic table.

• Atomic Number: Number of protons in the nucleus.

• Atomic Mass: Weighted average mass of an element's isotopes.

Chemical Formulas

Chemical formulas show the types and numbers of atoms in a compound.

• Example: represents water, with two hydrogen atoms and one oxygen atom.

Physical vs. Chemical Changes

Changes in matter can be physical or chemical.

• Physical Change: Alters form but not composition (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., burning wood).

Chemical Equations

Chemical equations represent chemical reactions, showing reactants and products.

• Reactants: Substances present before the reaction.

• Products: Substances formed by the reaction.

• Example:

Balancing Chemical Equations

Balanced equations have equal numbers of each atom on both sides.

• Rule: Adjust coefficients to balance atoms.

• Example:

Unit Conversions and Calculations

Converting between units is essential for solving chemistry problems.

• Example: Converting grams to kilograms:

Significant Figures

Significant figures reflect the precision of a measurement.

• Rule: All nonzero digits are significant; zeros between nonzero digits are significant.

• Example: 0.00450 has three significant figures.

Percent Calculations

Percent calculations are used to express ratios and concentrations.

• Formula:

Scientific Notation

Scientific notation expresses very large or small numbers concisely.

• Example: 0.00056 =

Mass, Weight, Volume, Density

These are fundamental physical properties.

• Mass: Amount of matter (unit: grams, g).

• Weight: Force of gravity on mass (unit: newtons, N).

• Volume: Space occupied (unit: liters, L).

• Density: Mass per unit volume (unit: g/mL or g/cm3).

• Formula:

Temperature Scales and Conversions

Three temperature scales are used in chemistry: Celsius, Fahrenheit, and Kelvin.

• Celsius (°C): Water freezes at 0°C, boils at 100°C.

• Fahrenheit (°F): Water freezes at 32°F, boils at 212°F.

• Kelvin (K): Absolute zero is 0 K; water freezes at 273.15 K.

• Conversions:

Energy: Potential and Kinetic

Energy is the capacity to do work. It exists as potential (stored) or kinetic (motion).

• Potential Energy: Stored energy due to position.

• Kinetic Energy: Energy of motion.

States of Matter

Matter exists as solid, liquid, or gas, each with distinct properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, no definite shape.

• Gas: No definite shape or volume.

Accuracy vs. Precision

Accuracy is how close a measurement is to the true value; precision is how reproducible measurements are.

• Example: Multiple measurements close together (precise), but not near the true value (not accurate).

Chapter 2: Atoms and Radioactivity

Subatomic Particles

Atoms are composed of electrons, protons, and neutrons.

• Electron: Negative charge, found outside nucleus.

• Proton: Positive charge, found in nucleus.

• Neutron: Neutral, found in nucleus.

Atomic Mass and Mass Number

Atomic mass is the weighted average of isotopes; mass number is the sum of protons and neutrons.

• Formula:

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: and are isotopes of carbon.

Radioactivity and Nuclear Radiation

Some isotopes are radioactive and undergo decay, emitting radiation.

• Radioisotope: An isotope that is radioactive.

• Radioactive Decay: Spontaneous emission of particles or energy.

Forms of Radiation

Three main types: alpha, beta, and gamma.

• Alpha Particle: ; low penetration.

• Beta Particle: ; moderate penetration.

• Gamma Ray: High-energy photon; high penetration.

Penetration of Radiation

Alpha particles are stopped by paper, beta by aluminum, gamma by lead.

• Additional info: Gamma rays are most dangerous due to deep penetration.

Half-Life

Half-life is the time required for half of a radioactive sample to decay.

• Formula: where is the number of half-lives.

Units of Radioactivity

Radioactivity is measured in curie (Ci) and becquerel (Bq).

• 1 Ci: disintegrations per second.

• 1 Bq: 1 disintegration per second.

Chapter 3: Compounds - How Elements Combine

Energy Levels and Electrons

Electrons occupy energy levels (shells) around the nucleus.

• Maximum Electrons: per level, where is the energy level.

Valence Electrons and Octet Rule

Valence electrons are in the outermost shell; the octet rule states atoms tend to have eight valence electrons.

• Noble Gases: Have a full octet, making them stable.

Ions: Anions and Cations

Ions are atoms or molecules with a net charge.

• Cation: Positive ion (loss of electrons).

• Anion: Negative ion (gain of electrons).

• Isoelectronic: Having the same electron configuration as another atom/ion.

Polyatomic Ions

Polyatomic ions are groups of atoms with a charge.

• Example: (nitrate).

Naming Ions and Compounds

Ionic compounds are named by cation then anion; covalent compounds use prefixes.

• Example: is sodium chloride; is carbon dioxide.

Ionic vs. Covalent Compounds

Ionic compounds form from metals and nonmetals; covalent compounds form from nonmetals.

• Ionic Bond: Transfer of electrons.

• Covalent Bond: Sharing of electrons.

Predicting Bond Polarity

Bond polarity depends on electronegativity differences.

• Nonpolar Covalent: Electrons shared equally.

• Polar Covalent: Electrons shared unequally.

• Electronegativity Trend: Increases across a period, decreases down a group.

Lewis Structures

Lewis structures show valence electrons as dots around atomic symbols.

• Example: has two lone pairs on oxygen.

Molecular Geometry (VSEPR)

VSEPR theory predicts molecular shapes based on electron pair repulsion.

• Example: is tetrahedral.

Bonding Electrons vs. Lone Pairs

Bonding electrons are shared between atoms; lone pairs are unshared.

• Single Bond: 2 electrons.

• Double Bond: 4 electrons.

• Triple Bond: 6 electrons.

Mass, Moles, and Avogadro's Number

Conversions between mass, moles, and number of particles use molar mass and Avogadro's number.

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

• Avogadro's Number: particles per mole.

• Formula:

HTML Table: Comparison of Bond Types

HTML Table: Temperature Scale Comparison

Additional info: Some content was inferred and expanded for completeness, such as example formulas, table entries, and explanations of concepts.