Chapter 1: Introduction—Matter, Energy, and Measurement

Chemistry: Definition and Scope

Chemistry is the study of matter, its properties, and the changes it undergoes. It is central to our understanding of many science-related fields, including biology, physics, medicine, and environmental science.

• Matter: Anything that has mass and occupies space.

• Properties of matter: Characteristics that define and distinguish substances.

• Chemical changes: Transformations that alter the composition of matter.

Classification of Matter

Matter can be classified by its physical state and composition.

States of Matter

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

• Liquid: Definite volume but no definite shape; particles are less tightly packed than in solids.

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

Classification by Composition

• Substance: Matter with distinct properties and a composition that does not vary from sample to sample.

  • Element: Cannot be decomposed into simpler substances; made of one kind of atom.

  • Compound: Can be decomposed into simpler substances; made of atoms of two or more elements.

• Mixture: Combination of two or more substances in which each retains its own properties.

  • Homogeneous mixture (solution): Uniform composition throughout.

  • Heterogeneous mixture: Composition varies throughout the sample.

Atoms and Molecules

• Atoms: The basic building blocks of matter; each element consists of a unique type of atom.

• Molecules: Groups of atoms bonded together, representing the smallest unit of a compound that retains its chemical properties.

Representing Elements

• Elements are represented by symbols (one or two letters, first capitalized).

• Some symbols are derived from Latin or Greek names.

Law of Constant Composition (Law of Definite Proportions)

• Compounds always have the same relative number of atoms of each element in any sample.

Properties of Matter

Properties are used to identify and classify substances.

• Physical properties: Can be observed without changing the substance (e.g., color, odor, density, melting point, boiling point, hardness).

• Chemical properties: Can only be observed when a substance is changed into another substance (e.g., flammability, reactivity).

Intensive vs. Extensive Properties

• Intensive properties: Independent of the amount of substance (e.g., density, boiling point, color).

• Extensive properties: Dependent on the amount of substance (e.g., mass, volume, energy).

Physical and Chemical Changes

• Physical changes: Do not alter the composition of a substance (e.g., changes of state, temperature, volume).

• Chemical changes (reactions): Result in the formation of new substances (e.g., combustion, oxidation, decomposition).

Separation of Mixtures

Mixtures can be separated by exploiting differences in physical properties.

• Filtration: Separates solids from liquids.

• Distillation: Separates components based on differences in boiling points.

• Chromatography: Separates substances based on their ability to adhere to a solid surface.

Energy and Its Forms

• Energy: The capacity to do work or transfer heat.

• Work: Energy transferred when a force moves an object.

• Heat: Energy used to increase the temperature of an object.

• Kinetic energy: Energy of motion, depends on mass and velocity.

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

Units of Measurement

Quantitative measurements are essential in chemistry. The International System of Units (SI) is used for consistency.

• SI Base Units:

  • Mass: kilogram (kg)

  • Length: meter (m)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount of substance: mole (mol)

  • Volume: derived unit (liter, L; milliliter, mL)

• Metric Prefixes: Used to express multiples or fractions of base units (e.g., kilo-, centi-, milli-).

Temperature Scales

• Celsius (°C): 0°C is the freezing point, 100°C is the boiling point of water.

• Kelvin (K): Absolute temperature scale; 0 K is absolute zero.

• Conversion:

• Fahrenheit (°F): Not commonly used in science.

Volume Units

• Volume is a derived unit: ,

Density

• Density ():

• Common units: or

Energy Units

• Joule (J): SI unit of energy.

• calorie (cal):

• Nutritional Calorie (Cal):

Uncertainty in Measurement

• All measurements have some degree of uncertainty due to limitations of instruments and human error.

• Exact numbers: Known with complete certainty (e.g., counting, definitions).

• Inexact numbers: Obtained by measurement; subject to uncertainty.

Precision vs. Accuracy

• Precision: How closely repeated measurements agree with each other.

• Accuracy: How closely a measurement agrees with the true value.

Significant Figures

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

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

Significant Figures in Calculations

• For addition/subtraction: Round to the least significant decimal place.

• For multiplication/division: Round to the same number of significant figures as the measurement with the fewest significant figures.

Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Set up ratios so that units cancel appropriately.

• Multiple conversions can be chained together as long as each is an equality.

• Example: To convert inches to centimeters, use .