Introduction to Chemistry

What is Chemistry?

Chemistry is the scientific discipline that seeks to understand the properties and behavior of matter by studying the behavior of atoms and molecules. It is foundational to many scientific fields and has broad applications in industry, health, engineering, and daily life.

• Matter: Anything that occupies space and has mass.

• Atoms: The fundamental building blocks of matter.

• Molecules: Groups of atoms bonded together.

• There are about 100 basic substances called elements that compose all matter.

Additional info: Advances in microscopy have allowed scientists to image atoms and molecules, providing direct evidence of their existence.

Classification of Matter

States of Matter

Matter can be classified according to its physical state: solid, liquid, or gas. Each state has distinct characteristics based on the arrangement and movement of particles.

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

• Liquid: Definite volume but no definite shape; particles are close but can move past one another.

• Gas: Neither definite shape nor volume; particles are far apart and move freely.

Composition of Matter

Matter can also be classified by its composition:

• Pure Substance: Composed of only one type of atom or molecule.

• Element: A pure substance that cannot be broken down into simpler substances.

• Compound: A pure substance composed of two or more elements in fixed, definite proportions.

• Mixture: Composed of two or more different types of atoms or molecules combined in variable proportions.

• Homogeneous Mixture (Solution): Uniform composition throughout; only one phase is observed.

• Heterogeneous Mixture: Non-uniform composition; more than one phase is observed.

Physical and Chemical Changes and Properties

Physical vs. Chemical Properties

• Physical Properties: Observed without changing the composition of the substance (e.g., boiling point, melting point, color, density).

• Chemical Properties: Observed only when a substance undergoes a chemical change (e.g., flammability, corrosiveness).

Physical vs. Chemical Changes

• Physical Change: The structure of matter is preserved (e.g., phase changes such as melting or boiling).

• Chemical Change: The structure of matter is changed (e.g., chemical reactions such as combustion).

Examples:

• Melting of a glacier: Physical change

• Human metabolism: Chemical change

• Burning Olympic torch: Chemical change

• Sodium chloride dissolution in water: Chemical change

• Snow formation: Physical change

• Leaves changing color: Physical change

Conservation Laws

Law of Conservation of Mass

Matter is neither created nor destroyed in a chemical reaction. During physical and chemical changes, the total amount of matter remains constant, although it may be converted from one type to another.

Law of Conservation of Energy

Energy is neither created nor destroyed. Chemical and physical changes are always accompanied by energy transfer, which can be in the form of heat, light, or work.

Forms of Energy

• Potential Energy: Due to position or composition.

• Kinetic Energy: Due to motion.

• Electrical Energy: Due to flow of electrical charges.

• Thermal Energy: Due to motion of atoms and molecules.

• Chemical Energy: A type of potential energy stored in chemical bonds.

Units of Measurement for Physical and Chemical Change

International System of Units (SI)

The SI system is the standard for scientific measurements. Common units include kilogram (kg), meter (m), second (s), mole (mol), and liter (L).

• Derived Units: Combinations of base units (e.g., m/s for speed, g/cm3 for density).

Dimensional Analysis and Unit Conversion

Dimensional analysis is a method for converting between units using conversion factors. For example, to convert 2.0 nm to μm:

When converting units raised to a power, the conversion factor must also be raised to that power. For example, converting cm3 to in3:

Volume and Density

• Volume: The amount of space an object occupies. Common units: m3, cm3, L.

• Density: The ratio of mass to volume. It is an intensive property, independent of the amount of substance.

Usual units: g/cm3 or g/mL.

Uncertainty in Measurement and Significant Figures

Uncertainty in Measured Quantities

All measurements have some degree of uncertainty, indicated by the last reported digit. The uncertainty affects the reliability and interpretation of scientific data.

Significant Figures

• All nonzero digits are significant.

• Interior zeros (between nonzero digits) are significant.

• Trailing zeros after a decimal point are significant.

• Leading zeros are not significant.

• Trailing zeros before an implied decimal point are ambiguous and should be avoided.

Examples:

• 0.0035 (2 SFs)

• 1.080 (4 SFs)

• 2371 (4 SFs)

• 2.9 × 105 (2 SFs)

• 100.00 (5 SFs)

Significant Figures in Calculations

• Multiplication/Division: The result has the same number of significant figures as the factor with the fewest significant figures.

• Addition/Subtraction: The result has the same number of decimal places as the quantity with the fewest decimal places.

Accuracy and Precision

• Accuracy: How close a measured value is to the true value.

• Precision: How reproducible a series of measurements are.