Ch.1 – Introduction: Matter, Energy, and Measurement (General Chemistry Study Notes)

Introduction to Chemistry: Matter, Energy, and Measurement

Scientific Method

The scientific method is a systematic approach to research and discovery in science. It involves making observations, forming hypotheses, conducting experiments, and drawing conclusions based on evidence.

• Observation: Gathering information from the natural world using the senses or instruments.

• Hypothesis: A tentative explanation or prediction that can be tested by experiments.

• Experiment: A controlled procedure to test the hypothesis.

• Data Analysis: Interpreting the results of experiments.

• Conclusion: Determining whether the hypothesis is supported or refuted.

• Theory: A well-substantiated explanation based on repeated testing and evidence.

• Law: A statement describing consistent natural phenomena.

Example: "Gravity is the reason that an object falls toward the Earth when released." This is a law.

Classification of Matter

Matter is anything that occupies space and has mass. It can be classified by its physical state and composition.

• Pure Substances: Matter with a fixed composition (elements and compounds).

• Mixtures: Physical combinations of two or more substances (homogeneous or heterogeneous).

Type	Description	Examples
Element	Cannot be broken down into simpler substances	Oxygen (O2), Gold (Au)
Compound	Composed of two or more elements chemically combined	Water (H2O), Sodium chloride (NaCl)
Homogeneous Mixture	Uniform composition throughout	Saltwater, Air
Heterogeneous Mixture	Non-uniform composition	Salad, Concrete

Example: Saltwater is a homogeneous mixture; salad is a heterogeneous mixture.

Physical and Chemical Changes

Changes in matter can be classified as physical or chemical:

• Physical Change: Alters the form or appearance but not the composition (e.g., melting ice, dissolving sugar in water).

• Chemical Change: Alters the composition, producing new substances (e.g., burning wood, rusting iron).

Example: Dissolving sugar in water is a physical change; burning wood is a chemical change.

Reversible and Irreversible Changes

• Reversible Change: Can be undone (e.g., melting/freezing water).

• Irreversible Change: Cannot be undone (e.g., burning paper).

Example: Dissolving sugar in water is reversible; burning paper is irreversible.

Chemical and Physical Properties

• Chemical Property: Observed during a chemical reaction (e.g., flammability, reactivity with acids).

• Physical Property: Can be measured without changing the substance (e.g., color, melting point, density).

Example: Mercury is a silvery liquid at 25°C (physical property); sodium reacts violently with water (chemical property).

Nature of Energy

Energy is the capacity to do work or produce heat. In chemistry, energy is associated with physical and chemical changes.

• Kinetic Energy: Energy of motion ($KE = \frac{1}{2}mv^2$)

• Potential Energy: Stored energy due to position ($PE = mgh$)

Type of Energy	Example
Kinetic	Moving car, flowing water
Potential	Stretched spring, water behind a dam

Energy Conversion Factors: 1 calorie (cal) = 4.184 joules (J); 1 kilowatt-hour (kWh) = 3.60 × 106 J

Scientific Notation

Used to express very large or small numbers conveniently.

• Format: $a \times 10^n$ where 1 ≤ a < 10 and n is an integer.

• Example: 6,800,000 = $6.8 \times 10^6$

SI Units and Measurements

The International System of Units (SI) is the standard for scientific measurements.

Physical Quantity	Unit Name	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol

Area: $A = l \times w$ Volume: $V = l \times w \times h$

Metric Prefixes

Metric prefixes indicate multiples or fractions of base units.

Prefix	Symbol	Multiplier
kilo	k	$10^3$
centi	c	$10^{-2}$
milli	m	$10^{-3}$
micro	μ	$10^{-6}$
nano	n	$10^{-9}$

Temperature

• Kelvin (K): SI unit for temperature.

• Celsius (°C): Commonly used in chemistry.

• Fahrenheit (°F): Used in the United States.

Temperature conversions:

• $K = °C + 273.15$

• $°C = \frac{5}{9}(°F - 32)$

• $°F = \frac{9}{5}°C + 32$

Significant Figures

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

• Rules for counting 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.

• When multiplying/dividing, the result has the same number of significant figures as the measurement with the fewest significant figures.

• When adding/subtracting, the result has the same number of decimal places as the measurement with the fewest decimal places.

Conversion Factors and Dimensional Analysis

Conversion factors are ratios used to express a quantity in different units. Dimensional analysis is a method for converting between units using conversion factors.

• General strategy: Multiply the given amount by the conversion factor so that units cancel appropriately.

Example: To convert 115 minutes to hours: $115 \text{ min} \times \frac{1 \text{ hr}}{60 \text{ min}} = 1.92 \text{ hr}$

Density

Density is the amount of mass per unit volume.

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

• Units: g/cm3 for solids and liquids; g/L for gases

Example: If a metal has a mass of 21.4 g and a volume of 10.0 cm3, its density is $2.14 \text{ g/cm}^3$.

Density of Geometric and Non-Geometric Objects

• For regular shapes, use geometric formulas for volume (cube: $V = a^3$, sphere: $V = \frac{4}{3}\pi r^3$).

• For irregular shapes, use water displacement to find volume.

Example: If a solid raises the water level from 200 mL to 265 mL, its volume is 65 mL.

*Additional info: Some examples and explanations have been expanded for clarity and completeness. Practice problems and diagrams referenced in the original notes are not reproduced here but are recommended for further study.*