Introduction to Chemistry: Exam 1 Study Guide

Chapter 1: Foundations of Chemistry

The Scientific Method

The scientific method is a systematic approach used in scientific study. It involves making observations, forming hypotheses, conducting experiments, and drawing conclusions.

• Observation: Gathering information using the senses.

• Hypothesis: A testable explanation for an observation.

• Experiment: A procedure to test the hypothesis.

• Theory: A well-substantiated explanation of some aspect of the natural world.

• Law: A statement that describes an observable occurrence in nature that appears to always be true.

Example: The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction.

Chapter 2: Measurement and Problem Solving in Chemistry

Scientific Notation

Scientific notation is a way to express very large or very small numbers using powers of ten.

• General form: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Example: $6,500 = 6.5 \times 10^3$

Significant Figures

Significant figures are the digits in a measurement that are known with certainty plus one estimated digit.

• Rules for identifying significant figures in a number.

• Rules for calculations:

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

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

• Example: $2.34 \times 1.2 = 2.8$ (rounded to two significant figures)

Units of Measurement and Conversions

Measurements in chemistry use the SI (International System of Units). Common units include meter (m), kilogram (kg), second (s), and mole (mol).

• Unit conversions use conversion factors to change from one unit to another.

• Example: $1 \text{ L} = 1000 \text{ cm}^3$

Dimensional Analysis

Dimensional analysis is a method for solving problems using conversion factors to move from one unit to another.

• Set up the problem so that units cancel appropriately.

• Example: Convert 5.0 g to mg: $5.0 \text{ g} \times \frac{1000 \text{ mg}}{1 \text{ g}} = 5000 \text{ mg}$

Density

Density is the mass per unit volume of a substance.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is $5 \text{ g/cm}^3$.

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified by its physical state and composition.

• States of matter: Solid, liquid, gas.

• Composition:

  • Elements: Pure substances made of one type of atom.

  • Compounds: Substances made from two or more elements chemically combined.

  • Mixtures: Physical combinations of two or more substances.

Physical and Chemical Properties

Physical properties 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).

Conservation of Mass and Energy

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Kinetic Energy and Heat

Kinetic energy is the energy of motion. Heat is the transfer of energy due to temperature difference.

• Formula for kinetic energy: $KE = \frac{1}{2}mv^2$

Units of Energy

• Joule (J): SI unit of energy.

• Calorie (cal): 1 cal = 4.184 J.

• Calorie (Cal): Used in food, 1 Cal = 1000 cal.

Chapter 4: Atomic Theory and the Periodic Table

Dalton’s Atomic Theory

John Dalton proposed that matter is made of indivisible particles called atoms, which combine in simple whole-number ratios to form compounds.

The Nuclear Atom

The atom consists of a nucleus containing protons and neutrons, with electrons orbiting the nucleus.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle orbiting the nucleus.

Properties of Subatomic Particles

Particle	Charge	Location	Relative Mass
Proton	+1	Nucleus	1
Neutron	0	Nucleus	1
Electron	-1	Outside nucleus	~0

The Periodic Table and Trends

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

• Trends: Atomic radius, ionization energy, and electronegativity change predictably across periods and groups.

Ions and Ionic Formation

Ions are formed when atoms gain or lose electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Isotopes and Atomic Mass

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

• Number of neutrons = mass number - atomic number.

• Average atomic mass is calculated using the relative abundance and mass of each isotope:

$\text{Average atomic mass} = \sum (\text{fractional abundance} \times \text{isotopic mass})$

Example: If an element has two isotopes with masses 10.0 amu (20%) and 11.0 amu (80%), average atomic mass = $(0.20 \times 10.0) + (0.80 \times 11.0) = 10.8$ amu.