Chapter 1: Chemistry in Our Lives

Definitions and Basic Concepts

• Chemistry: The study of matter, its properties, the changes it undergoes, and the energy associated with those changes.

• Chemicals: Substances with a definite composition; also called chemical substances.

• Matter: Anything that has mass and occupies space.

• Identifying Chemicals and Matter: All physical objects are matter; chemicals are specific types of matter with defined composition (e.g., water, sodium chloride).

Significant Figures

• Significant Figures: Digits in a number that are known with certainty plus one estimated digit.

• Determining Significant Figures:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Calculations with Significant Figures:

  • 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 the final answer to the correct number of significant figures.

Scientific Notation

• Scientific Notation: Expresses numbers as a product of a coefficient (between 1 and 10) and a power of 10.

• Example:

• Standard Notation: Converting from scientific notation to decimal form. Example:

Chapter 2: Chemistry and Measurements

Units of Measurement

• Base SI Units:

  • Mass: kilogram (kg)

  • Length: meter (m)

  • Volume: liter (L) or cubic meter (m3)

  • Time: second (s)

  • Temperature: kelvin (K), degrees Celsius (°C)

• Metric Prefixes: Used to express multiples or fractions of units.

  • Mega (M):

  • Kilo (k):

  • Centi (c):

  • Milli (m):

  • Micro (\mu):

  • Nano (n):

• Unit Conversions: Use conversion factors based on metric prefixes to convert between units (e.g., meters to centimeters).

• Percentage as a Conversion Factor: Percentages can be used to relate parts to whole in calculations.

Density

• Definition: Density is mass per unit volume.

• Formula:

• Units: Commonly expressed in g/mL or g/cm3.

• Calculations:

  • To find density:

  • To find mass:

  • To find volume:

Chapter 3: Matter and Energy

States of Matter

• Solid (s): Definite shape and volume; particles are closely packed in a fixed arrangement.

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

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

Classification of Matter

• Pure Substance:

  • Element: Cannot be broken down into simpler substances (e.g., O2, Fe).

  • Compound: Composed of two or more elements chemically combined (e.g., H2O, NaCl).

• Mixture:

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

  • Heterogeneous: Non-uniform composition (e.g., salad, sand in water).

Physical and Chemical Properties/Changes

• Physical Properties: Characteristics observed without changing the substance's identity (e.g., color, melting point).

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

• Physical Change: Change in state or appearance without changing composition (e.g., melting ice).

• Chemical Change: Substance is transformed into a different substance (e.g., rusting iron).

Energy

• Definition: The capacity to do work or produce heat.

• SI Unit: Joule (J)

• Kinetic Energy: Energy of motion.

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

Nutritional Calorie

• Calorie (cal): Amount of energy needed to raise the temperature of 1 g of water by 1°C.

• kilocalorie (kcal): 1 kcal = 1000 cal.

• Nutritional Calorie (Cal): 1 Cal = 1 kcal = 1000 cal.

• Conversions:

• Energy Values of Food Types:

  • Carbohydrates: 4 cal/g

  • Proteins: 4 cal/g

  • Fats: 9 cal/g

Chapter 4: Atoms and Elements

Atomic Structure

• Subatomic Particles:

  • Proton: Positive charge (+1), mass ≈ 1 amu, located in nucleus.

  • Neutron: No charge (0), mass ≈ 1 amu, located in nucleus.

  • Electron: Negative charge (–1), mass ≈ 0.0005 amu, located outside nucleus.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons ().

• Number of Neutrons:

• Number of Electrons: For neutral atoms, equal to number of protons.

Isotopes

• Definition: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Atomic Mass of Isotopes: Weighted average based on natural abundance.

  • Formula:

The Periodic Table

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

• Periods: Horizontal rows.

• Groups: Vertical columns.

• Classes of Elements:

  • Metals: Good conductors, malleable, ductile, shiny.

  • Metalloids: Properties intermediate between metals and nonmetals.

  • Nonmetals: Poor conductors, brittle, dull.

  • Special Groups:

    • Halogens: Group 17

    • Alkali Metals: Group 1

    • Noble Gases: Group 18

• Trends in Periodic Properties:

  • Atomic Size: Generally increases down a group and decreases across a period.

Ions

• Definition: Atoms or molecules that have gained or lost electrons, acquiring a charge.

• Charge Calculation:

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

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

• Determining Subatomic Particles in Ions: Adjust electron count based on charge.

• Valence Electrons: Number of electrons in the outermost shell; for main group elements, equals group number.

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

Chapter 6: Ionic and Molecular Compounds

Lewis Symbols

• Lewis Symbol: Element symbol surrounded by dots representing valence electrons.

• Number of Dots: Equals the number of valence electrons.

Ionic Bonds

• Definition: Chemical bond formed by the transfer of electrons from a metal to a nonmetal.

• Representation: Use Lewis symbols to show electron transfer.

• Identifying Ionic Compounds: Typically formed between metals and nonmetals (e.g., NaCl).

Covalent Bonds

• Definition: Chemical bond formed by the sharing of electrons between nonmetals.

• Representation: Use Lewis symbols to show shared electron pairs.

• Identifying Covalent Compounds: Typically formed between nonmetals (e.g., H2O, CO2).

Writing Formulas and Naming Compounds

• Formulas:

  • Monoatomic Ions: Single atom with a charge (e.g., Na+, Cl–).

  • Polyatomic Ions: Group of atoms with a charge (e.g., SO42–).

• Naming Ionic Compounds: Name the cation first, then the anion (e.g., sodium chloride).

• Naming Covalent Compounds: Use prefixes to indicate number of atoms (e.g., carbon dioxide, dinitrogen monoxide).

Properties of Ionic and Covalent Compounds

• Ionic Compounds: High melting/boiling points, conduct electricity when dissolved in water, usually solid at room temperature.

• Covalent Compounds: Lower melting/boiling points, do not conduct electricity, can be solid, liquid, or gas at room temperature.

Electrolytes and Non-Electrolytes

• Electrolytes: Substances that conduct electricity when dissolved in water (typically ionic compounds).

• Non-Electrolytes: Substances that do not conduct electricity in solution (typically covalent compounds).

Additional info: This study guide covers foundational topics in general chemistry, including measurement, atomic structure, periodic trends, chemical bonding, and basic nomenclature. Mastery of these concepts is essential for success in subsequent chemistry topics.