GOB Chemistry Study Guide: Measurement, Matter, Atomic Structure, and Chemical Bonding

Chapter 1: Introduction and Measurement

Definition of Chemistry and Matter

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. Matter is anything that has mass and occupies space.

• Identify chemicals and matter: Recognize substances as chemicals if they have a definite composition and properties.

Significant Figures and Scientific Notation

Significant figures are the digits in a measurement that are known with certainty plus one estimated digit. Scientific notation expresses very large or very small numbers in the form $a \times 10^n$.

• Express numbers to the correct significant figures: Use the correct number of digits based on measurement precision.

• Determine number of significant figures: Count all nonzero digits and any zeros between them or after a decimal point.

• Perform calculations and consider significant figures appropriately:

  • Multiplication/Division: The result should have as many significant figures as the measurement with the fewest significant figures.

  • Addition/Subtraction: The result should have as many decimal places as the measurement with the fewest decimal places.

• Rounding: Round to the correct number of significant figures as required by the calculation.

• Use scientific notation to express standard notation: For example, $0.015$ is written as $1.5 \times 10^{-2}$.

• Use standard notation to express scientific notation: For example, $1.5 \times 10^{-2}$ is $0.015$.

Chapter 2: Chemistry and Measurements

Units of Measurement

Measurements in chemistry require standard units for mass, length, volume, time, and temperature. The metric system is commonly used.

• Identify standards and units of measurement: Common metric units include meter (m), gram (g), liter (L), and second (s).

• Identify metric units and prefixes: Examples include kilo- (k, $10^3$), centi- (c, $10^{-2}$), milli- (m, $10^{-3}$), micro- ($\mu$, $10^{-6}$).

• Convert units using prefixes: For example, convert meters to centimeters, liters to milliliters, etc.

• Use percentages as a conversion factor: Percentages can be used to relate parts to a whole in calculations.

Density

Density is a physical property defined as mass per unit volume. It is used to identify substances and convert between mass and volume.

• Definition: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Calculations:

  • Determine density from mass and volume.

  • Determine mass from density and volume.

  • Determine volume from density and mass.

Chapter 3: Matter and Energy

States of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has characteristic properties based on the arrangement and movement of particles.

• Solid: Definite shape and volume; particles are closely packed and vibrate in place.

• Liquid: Definite volume but no definite shape; particles are close but can move past each other.

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

Classification of Matter

Matter can be classified as pure substances or 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 (uniform) or heterogeneous (non-uniform).

Chemical and Physical Properties and Changes

• Physical properties: Characteristics that can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical properties: Describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical change: Alters the form but not the identity of a substance (e.g., melting, boiling).

• Chemical change: Produces new substances with different properties (e.g., rusting, burning).

Energy in Chemistry

Energy is the capacity to do work or produce heat. It exists in various forms, including kinetic and potential energy.

• Kinetic energy: Energy of motion.

• Potential energy: Stored energy due to position or composition.

• Units: Energy is measured in calories (cal) and joules (J).

• Calorie: 1 cal = 4.184 J.

• Food energy: Measured in kilocalories (kcal); 1 kcal = 1000 cal.

Chapter 4: Atoms and Elements

Atomic Structure

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Number of neutrons: $\text{Number of neutrons} = \text{Mass number} - \text{Atomic number}$

• Number of electrons: In a neutral atom, equal to the number of protons.

Isotopes

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

• Definition: Atoms with the same atomic number but different mass numbers.

• Isotopic abundance: The relative amount of each isotope in a natural sample of the element.

The Periodic Table

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

• Periodic Law: The properties of elements are periodic functions of their atomic numbers.

• Periods: Horizontal rows.

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

• Classification: Elements are classified as metals, nonmetals, and metalloids.

• Special groups: Alkali metals, alkaline earth metals, halogens, noble gases.

• Trends: Atomic size, metallic character, and ionization energy vary predictably across periods and groups.

Ions and Valence Electrons

Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge. Valence electrons are the outermost electrons involved in chemical bonding.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

• Valence electrons: Number of valence electrons = group number for main group elements.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons.

Chapter 6: Ionic and Molecular Compounds

Lewis Symbols and Ionic Bonding

Lewis symbols represent valence electrons as dots around the element symbol. Ionic bonds form when electrons are transferred from metals to nonmetals, resulting in oppositely charged ions.

• Draw Lewis symbols: Show valence electrons for main group elements.

• Ionic bond: Electrostatic attraction between cations and anions.

• Use Lewis symbols to represent ionic bonding: Show electron transfer and resulting ions.

Covalent Bonding and Molecular Compounds

Covalent bonds form when two nonmetals share electrons. Molecular compounds are composed of molecules held together by covalent bonds.

• Draw Lewis structures: Show shared pairs of electrons between atoms.

• Write formulas and name compounds: Use systematic naming for binary and polyatomic compounds.

• Name ionic and covalent compounds: Follow IUPAC rules for naming, including transition metals and polyatomic ions.

Electrolytes and Non-Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water; non-electrolytes do not.

• Define electrolytes: Ionic compounds that dissociate into ions in solution.

• Define non-electrolytes: Molecular compounds that do not produce ions in solution.

Example Table: Classification of Matter

Type	Definition	Examples
Element	Pure substance made of one type of atom	O2, Fe, H2
Compound	Pure substance made of two or more elements chemically combined	H2O, CO2
Homogeneous Mixture	Uniform composition throughout	Salt water, air
Heterogeneous Mixture	Non-uniform composition	Salad, sand and water

Additional info: Some explanations and examples have been expanded for clarity and completeness based on standard GOB Chemistry curriculum.