General Chemistry: Chapter 1 – Matter, Measurement, and Problem Solving

General Chemistry: Chapter 1 – Matter, Measurement, and Problem Solving

Chemistry and Matter

Chemistry is the scientific study of matter, its properties, and the changes it undergoes. Matter is composed of atoms and molecules, and chemistry seeks to understand the behavior of matter by examining the behavior of these fundamental particles.

• Matter: Anything that occupies space and has mass.

• Atoms: The submicroscopic particles that are the building blocks of ordinary matter.

• Molecules: Groups of atoms bonded together in specific geometrical arrangements.

The Classification of Matter

Matter can be classified by its physical state and its composition. The physical states are solid, liquid, and gas, while the composition can be elements, compounds, or mixtures.

• States of Matter:

  • Solids: Fixed volume and rigid shape (crystalline or amorphous).

  • Liquids: Fixed volume but not a fixed shape.

  • Gases: Compressible and take the shape and volume of their container.

• Classification by Components:

  • Elements: Pure substances consisting of only one type of atom.

  • Compounds: Substances composed of two or more types of atoms chemically bonded.

  • Mixtures: Physical combinations of two or more substances.

Properties of Matter

Properties of matter are divided into intensive and extensive properties.

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

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

Physical and Chemical Properties

Physical properties can be observed without changing the composition of a substance, while chemical properties are observed only when a substance undergoes a chemical change.

• Physical Properties: Odor, taste, color, appearance, melting point, boiling point, density.

• Chemical Properties: Flammability, corrosiveness, acidity, toxicity.

Physical and Chemical Changes

Changes in matter are classified as physical or chemical.

• Physical Change: Alters only the state or appearance, not the composition. Example: Boiling water changes liquid to gas, but the molecules remain H2O.

• Chemical Change: Alters the composition of matter. Atoms rearrange to form new substances. Example: Rusting of iron.

Example Classification

• Natural gas burns in a stove – Chemical change

• Liquid propane evaporates – Physical change

• Liquid propane burns in a flame – Chemical change

• Bicycle frame rusts – Chemical change

Energy in Physical and Chemical Change

Energy is the capacity to do work, and work is the action of a force through a distance. Energy is fundamental to both physical and chemical changes.

• Kinetic Energy: Energy associated with motion.

• Thermal Energy: A type of kinetic energy related to temperature.

• Potential Energy: Energy associated with position or composition (e.g., bonding energy).

Measurement

Measurement in chemistry requires accuracy and precision.

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

• Precision: How close a series of measurements are to one another.

Significant Figures

Significant figures are important in reporting measurements and calculations.

• All nonzero digits are significant.

• Interior zeroes (between nonzero digits) are significant.

• Leading zeroes are not significant.

• Trailing zeroes after a decimal point are significant.

• Trailing zeroes before a decimal point are ambiguous unless specified by scientific notation.

Rule	Example
All nonzero digits	28.03
Interior zeroes	7.6301
Leading zeroes	0.0540
Trailing zeroes after decimal	45.000
Trailing zeroes before decimal	1200 (ambiguous)

Significant Figure Rules for Calculations

• Multiplication and Division: Result has the same number of significant figures as the factor with the fewest significant figures. $1.052 imes 12.504 imes 0.53 = 6.7208 ightarrow 6.7$ (2 sig. figures)

• Addition and Subtraction: Result has the same number of decimal places as the quantity with the fewest decimal places. $5.679 - 0.221 = 5.458 ightarrow 5.7$ (1 decimal place)

The Units of Measurement

Chemistry uses standard units to specify measurements. The metric system and the International System of Units (SI) are most common.

Quantity	Unit	Symbol
Length	Meter	m
Mass	Kilogram	kg
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol
Electric current	Ampere	A
Luminous intensity	Candela	cd

The Kelvin: A Measure of Temperature

The Kelvin (K) is the SI unit of temperature. Temperature measures the average kinetic energy of the particles in matter and determines the direction of heat transfer.

• Kelvin scale: Assigns 0 K (absolute zero) to the coldest temperature possible.

• Absolute zero: $-273^ ext{o}C$ or $-459^ ext{o}F$; molecular motion virtually stops.

• Conversion: $K = ^ ext{o}C + 273.15$

Prefix Multipliers

SI units use prefix multipliers to indicate powers of ten.

Prefix	Symbol	Multiplier
kilo	k	1,000
centi	c	0.01
milli	m	0.001
micro	μ	0.000001
nano	n	0.000000001
pico	p	0.000000000001

Unit Conversion and Dimensional Analysis

Unit conversion is essential in chemistry calculations. Dimensional analysis uses units as a guide to solving problems.

• Conversion factor: Fractional quantity of a unit equation with the units being converted from on the bottom and the units being converted to on the top.

• General form: $\text{Information given} \times \text{conversion factor(s)} = \text{information sought}$

• Example: $100 \text{ m} \times (1 \text{ cm}/0.01 \text{ m}) = 10,000 \text{ cm}$

Units Raised to a Power

When converting units raised to a power, raise both the number and the unit to the power.

• Example: $2.54 \text{ cm} = 1 \text{ in}$ $(2.54 \text{ cm})^2 = (1 \text{ in})^2$ $6.45 \text{ cm}^2 = 1 \text{ in}^2$

Examples: Volume and Density

• Volume Conversion: 1 gallon = 3.79 dm3, 1 dm = 10 cm $0.5 \text{ G} \times (3.79 \text{ dm}^3 / 1 \text{ G}) = 1.90 \text{ dm}^3$ $1.90 \text{ dm}^3 \times (1000 \text{ cm}^3 / 1 \text{ dm}^3) = 1900 \text{ cm}^3$

• Density: $d = \frac{\text{mass}}{\text{volume}}$ Unit of density in SI: kg/m3 Example comparison: $10 \text{ kg/m}^3 = 1/1000 \text{ g/cm}^3$ $5 \text{ g/cm}^3$ is much denser than $10 \text{ kg/m}^3$

Additional info: These notes cover foundational concepts in general chemistry, including matter, measurement, and problem solving, and are suitable for introductory college-level study.