Chapter 1: Chemistry Basics - Matter and Measurement

Matter

Matter is anything that has mass and occupies space. Understanding matter is fundamental to chemistry, as it forms the basis for all chemical substances and reactions.

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

• Examples: Water, air, rocks, and living organisms.

Classification of Matter

Matter can be classified using a flow chart into pure substances and mixtures, which are further subdivided.

• 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

Elements are pure substances consisting of only one type of atom. Atoms are the smallest unit of an element that retains its properties.

• Element: Cannot be broken down by chemical means.

• Atom: Basic unit of matter; consists of protons, neutrons, and electrons.

Periodic Table of Elements

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

• Groups: Vertical columns; elements in a group share similar chemical properties.

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

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

• Metals: Groups 1A and 2A; typically shiny, conductive, and malleable.

• Halogens: Group 7A; highly reactive nonmetals.

• Noble Gases: Group 8A; 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 represent the elements and their ratios in a compound.

• Example: for water.

Physical vs. Chemical Changes

Physical changes alter the form of matter but not its chemical identity, while chemical changes produce new substances.

• Physical Change: Melting, boiling, dissolving.

• Chemical Change: Burning, rusting, reacting.

Chemical Equations

Chemical equations show reactants and products in a chemical reaction.

• Reactants: Starting substances.

• Products: Substances formed.

• Example:

Balancing Chemical Equations

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

• Method: Adjust coefficients to balance atoms.

Unit Conversions and Calculations

Converting between units is essential for quantitative chemistry.

• Example: Converting grams to kilograms.

Significant Figures

Significant figures reflect the precision of a measurement.

• Rule: All nonzero digits are significant; zeros may or may not be, depending on their position.

Percent Calculations

Percent calculations are used to express ratios and concentrations.

• Formula:

Scientific Notation

Scientific notation expresses very large or small numbers concisely.

• Example:

Mass, Weight, Volume, Density

These are fundamental physical properties.

• Mass: Amount of matter (grams, kg).

• Weight: Force due to gravity (newtons).

• Volume: Space occupied (liters, mL).

• Density:

Density Calculations

Density is calculated by dividing mass by volume.

• Formula:

Temperature Scales and Conversions

Three temperature scales are used: Celsius, Fahrenheit, Kelvin.

• Conversion Formulas:

Energy: Potential and Kinetic

Energy is the capacity to do work; it exists as potential or kinetic energy.

• Potential Energy: Stored energy.

• 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, variable shape.

• Gas: Variable shape and 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 close to 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: No charge, 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

Radioactivity and Nuclear Radiation

Radioactive atoms undergo decay, emitting radiation.

• Alpha Particle: , 2 protons and 2 neutrons, charge +2.

• Beta Particle: , electron, charge -1.

• Gamma Ray: , electromagnetic radiation, no charge.

Penetration of Radiation

Different types of radiation penetrate matter to varying degrees.

• Alpha: Least penetrating; stopped by paper.

• Beta: Moderate; stopped by plastic or glass.

• Gamma: Most penetrating; requires lead or thick concrete.

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 Electron Configuration

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 electrons in their valence shell.

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

Ions: Anions and Cations

Ions are atoms or molecules with a net charge.

• Cation: Positive charge (loss of electrons).

• Anion: Negative charge (gain of electrons).

Isoelectronic Species

Isoelectronic species have the same number of electrons.

• Example: and

Polyatomic Ions

Polyatomic ions are groups of atoms with a charge.

• Example: (nitrate)

Naming Ions and Compounds

Ions and compounds are named according to specific rules.

• Ionic Compounds: Name cation first, then anion.

• Covalent Compounds: Use prefixes to indicate number of atoms.

Ionic and Covalent Bonds

Ionic bonds form between metals and nonmetals; covalent bonds form between nonmetals.

• Ionic Compound: Transfer of electrons.

• Covalent Compound: Sharing of electrons.

Predicting Compound Formulas

Use the criss-cross method to determine ionic compound formulas.

• Example: and form

Lewis Structures

Lewis structures show valence electrons as dots around atomic symbols.

• Bonding Electrons: Shared between atoms.

• Lone Pairs: Not shared.

Bond Types: Polar vs. Nonpolar

Bonds can be polar (unequal sharing) or nonpolar (equal sharing).

• Electronegativity: Determines bond polarity.

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

Molecular Geometry (VSEPR Theory)

VSEPR theory predicts the shape of molecules based on electron pair repulsion.

• Example: is bent; is linear.

Mass, Moles, and Avogadro's Number

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

• Avogadro's Number: particles per mole.

• Formula:

Molar Mass Calculations

Molar mass is the sum of atomic masses in a compound.

• Example: g/mol

Bonding Electrons and Lone Pairs

Bonding electrons are shared; lone pairs are not.

• Single Bond: 2 electrons.

• Double Bond: 4 electrons.

• Triple Bond: 6 electrons.

Summary Table: Types of Bonds and Their Properties

Electronegativity Trends

Electronegativity increases across a period and decreases down a group.

• Highest: Fluorine

• Lowest: Cesium

Additional info:

Some content was inferred and expanded for completeness, including examples, formulas, and explanations of concepts such as VSEPR theory, percent calculations, and the criss-cross method for ionic compounds.