Chapter 1: Introduction to Matter, Energy, and Measurement – General Chemistry Study Notes

Introduction to Chemistry

What is Chemistry?

Chemistry is the scientific study of matter, its properties, and the changes it undergoes. It is central to understanding many science-related fields and has several branches, including organic, inorganic, physical, analytical, and biochemistry.

• Matter: Anything that has mass and occupies space (e.g., people, water, air).

• Branches of Chemistry: Organic, inorganic, physical, analytical, biochemistry, and more.

Classification of Matter

Methods of Classification

Matter can be classified by its physical state and by its composition.

• State of Matter: Solid, liquid, gas.

• Composition of Matter: Element, compound, mixture.

The States of Matter

• Solid: Fixed shape and volume, not compressible (e.g., ice).

• Liquid: Fixed volume, takes shape of container, not compressible (e.g., water).

• Gas: No fixed shape or volume, compressible (e.g., water vapor).

Classification by Composition

Matter can be classified as pure substances or mixtures.

• Pure Substances: Elements and compounds.

• Mixtures: Homogeneous (solutions) and heterogeneous mixtures.

Substances

• Element: Cannot be decomposed into simpler substances. Each element is made of unique atoms.

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

Law of Constant Composition: Compounds have a definite composition; the relative number of atoms of each element is the same in any sample.

Mixtures

• Heterogeneous Mixture: Composition varies throughout the sample (e.g., sand in water).

• Homogeneous Mixture (Solution): Uniform composition throughout (e.g., salt water).

Summary Table: Classification of Matter

Type	Definition	Examples
Element	Cannot be broken down into simpler substances	O2, Fe
Compound	Composed of two or more elements in fixed ratios	H2O, CO2
Homogeneous Mixture	Uniform composition throughout	Salt water, air
Heterogeneous Mixture	Non-uniform composition	Salad, sand in water

Atomic and Molecular Perspective

The properties of matter are determined by the composition (elements) and structure (arrangement) of atoms and molecules.

Properties and Changes of Matter

Physical vs. Chemical Properties

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

• Chemical Properties: Describe how a substance changes into another (e.g., flammability, reactivity).

Intensive vs. Extensive Properties

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

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

Physical vs. Chemical Changes

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

• Chemical Changes: Result in new substances (e.g., combustion, oxidation).

Separation of Mixtures

Mixtures can be separated by exploiting differences in physical properties.

• Filtration: Separates solids from liquids.

• Distillation: Uses differences in boiling points to separate components.

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

Measurement in Chemistry

Numbers and Chemistry

• Many chemical concepts are quantitative and require measurement.

• Key concepts: units of measurement, uncertainty, significant figures, dimensional analysis.

SI Units (International System of Units)

Physical Quantity	Name of Unit	Abbreviation
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

Metric System and Prefixes

Prefixes are used to indicate multiples or fractions of units.

Prefix	Symbol	Factor
kilo	k	103
centi	c	10-2
milli	m	10-3
micro	μ	10-6
nano	n	10-9
pico	p	10-12

Derived Units

• Formed by combining SI base units (e.g., speed in m/s).

• Common derived units: volume (m3), density (kg/m3), energy (joule, J).

Volume

• SI unit: cubic meter (m3), but liter (L) is commonly used.

• 1 L = 1 dm3 = 1000 cm3

Temperature

• Measured in Celsius (°C) and Kelvin (K).

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

• Fahrenheit to Celsius: $^ ext{o}C = \frac{5}{9}(F - 32)$

• Celsius to Fahrenheit: $F = \frac{9}{5}^ ext{o}C + 32$

Energy

Forms of Energy

• Kinetic Energy: Energy of motion. $E_k = \frac{1}{2}mv^2$

• Potential Energy: Energy due to position or composition.

Energy is measured in joules (J). 1 Calorie (Cal) = 1,000 calories (cal).

Density

• Density is a physical property: $d = \frac{m}{V}$

• Common units: g/cm3 (solids, liquids), g/L (gases).

Measurement and Uncertainty

Numbers in Science

• Exact Numbers: Known with certainty (e.g., 1 dozen = 12).

• Inexact Numbers: Measured values, subject to uncertainty.

Uncertainty, Accuracy, and Precision

• Uncertainty: Last digit in a measurement is estimated.

• Precision: How closely repeated measurements agree.

• Accuracy: How close a measurement is to 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.

Scientific Notation

• Expresses numbers as $m \times 10^n$ where $1 \leq m < 10$ and $n$ is an integer.

• Example: 45,000 = $4.5 \times 10^4$

Significant Figures in Calculations

• Multiplication/Division: Result has as many significant figures as the factor with the fewest significant figures.

• Addition/Subtraction: Result has as many decimal places as the number with the fewest decimal places.

Dimensional Analysis

Dimensional analysis is used to convert units using conversion factors.

• Set up ratios so units cancel appropriately.

• Example: To convert 100 cm to inches, use $1 \text{ in} = 2.54 \text{ cm}$.

Best Practices

• Always include units in answers.

• Use leading zeros for numbers less than 1 (e.g., 0.25, not .25).

• Keep track of units to ensure answers make sense.