Chapter 2 – Chemistry and Measurements

Units of Measurement

In chemistry, measurements are essential for quantifying substances and describing physical properties. Standard units ensure consistency and accuracy in scientific communication.

• SI Units: The International System of Units (SI) is the standard for scientific measurements. Common SI units include the meter (m) for length, kilogram (kg) for mass, and second (s) for time.

• Metric System: A decimal-based system used globally in science. Prefixes such as kilo-, centi-, and milli- denote multiples or fractions of base units.

• Example: 1 kilometer (km) = 1000 meters (m).

Scientific Notation

Scientific notation expresses very large or very small numbers in a compact form, making calculations easier and reducing errors.

• Format: A number is written as a × 10n, where 1 ≤ a < 10 and n is an integer.

• Example: 0.00056 = 5.6 × 10-4

Significant Figures

Significant figures (sig figs) indicate the precision of a measured value. The number of significant figures reflects the certainty in a measurement.

• Measured vs. Exact Numbers: Measured numbers have limited precision; exact numbers (like counted items or defined quantities) have infinite precision.

• Rules:

  • 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.

• Rounding: When performing calculations, round the final answer to the correct number of significant figures based on the operation.

• Example: 0.00450 has three significant figures.

Conversion Factors

Conversion factors are ratios used to convert one unit to another without changing the value of a quantity.

• Setup: Multiply the original measurement by a fraction (conversion factor) so that units cancel appropriately.

• Example: To convert 5.0 cm to meters:

Density

Density is a physical property defined as mass per unit volume. It is useful for identifying substances and converting between mass and volume.

• Formula:

• Units: Commonly expressed in g/mL or g/cm3 for liquids and solids.

• Example: If a sample has a mass of 10.0 g and a volume of 2.0 mL, its density is

Chapter 3 – Matter and Energy

Classification of Matter

Matter is anything that has mass and occupies space. It can be classified based on its composition and uniformity.

• Pure Substances: Have a fixed composition and distinct properties.

  • Elements: Consist of only one type of atom (e.g., O2, Fe).

  • Compounds: Composed of two or more elements chemically combined in fixed ratios (e.g., H2O, NaCl).

• Mixtures: Physical combinations of two or more substances.

  • Homogeneous Mixtures (Solutions): Uniform composition throughout (e.g., saltwater).

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

States of Matter

Matter exists in different physical forms called states, each with distinct properties.

• 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.

Physical vs. Chemical Properties

Properties of matter can be classified as physical or chemical.

• Physical Properties: Can be observed or measured without changing the substance's identity (e.g., color, melting point, density).

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

Physical vs. Chemical Changes

Changes in matter can be physical or chemical.

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

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

Temperature

Temperature measures the average kinetic energy of particles in a substance. It is a fundamental parameter in chemistry.

• Common Scales: Celsius (°C), Kelvin (K), Fahrenheit (°F).

• Conversions:

Specific Heat

Specific heat is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

• Formula:

• Where:

  • = heat (J or cal)

  • = mass (g)

  • = specific heat (J/g·°C or cal/g·°C)

  • = change in temperature (°C)

• Example: Calculating the heat needed to warm 100 g of water by 10°C, with :

Energy and Nutrition

Energy is measured in joules (J) or calories (cal). In nutrition, the Calorie (with a capital C) is used, where 1 Calorie = 1 kilocalorie (kcal) = 1000 calories.

• Conversions:

• Application: Food labels use Calories to indicate energy content.

Summary Table: Classification of Matter

Summary Table: Temperature Conversions

Additional info:

• "Slash & burn" or "railroad tracks" refer to the method of setting up unit conversions using conversion factors in a stepwise, canceling-units approach.

• Mastering Chemistry and Pearson resources provide additional practice and tutorials for these foundational topics.