CHM 120 Chapter 1: Matter and Measurements – Structured Study Notes

Introduction to Chemistry

Chemistry as the Central Science

Chemistry is often referred to as the central science because it connects and overlaps with many other scientific disciplines, such as biology, physics, and environmental science.

• Chemistry is the study of matter and the changes it undergoes.

• Matter is anything that has mass and occupies space; it can be seen, touched, tasted, or smelled.

Elements and Symbols

Organization and Representation

Elements are the fundamental substances that make up all matter. They are organized in the periodic table and represented by unique chemical symbols.

• Periodic Table: A systematic arrangement of elements based on their properties.

• Chemical Symbols: Each element is represented by a one- or two-letter symbol. The first letter is always capitalized, and the second letter (if present) is lowercase (e.g., H for hydrogen, He for helium).

• Chemical Formulas: Indicate the number of atoms of each element in a compound using subscripts (e.g., H2O).

Quantities – Numbers and Units

Physical Quantities and Units

Measurements in chemistry require both a number and a unit to be meaningful. Units are essential for clarity and accuracy.

• Physical Quantity: A measurable property (e.g., mass, length, time).

• SI Units: The International System of Units is used for scientific measurements.

Physical Quantity	SI Unit	Abbreviation	Common Unit	Abbreviation
Mass	kilogram	kg	gram	g
Length	meter	m
Time	second	s
Temperature	kelvin	K	celsius	°C
Amount	mole	mol
Energy	joule	J	calorie	cal
Pressure	pascal	Pa	atmosphere	atm

• Metric Prefixes: Used to express larger or smaller quantities (e.g., kilo-, milli-, micro-).

• Scientific Notation: Used for very large or very small numbers. Example: $1.2 \times 10^6$.

Accuracy vs Precision

Definitions and Differences

Accuracy and precision are important concepts in measurement.

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

• Precision: How reproducible or consistent repeated measurements are.

• Example: If a balance consistently reads 5.0 g for a 5.0 g object, it is both accurate and precise.

Significant Figures

Rules and Calculations

Significant figures (sig figs) indicate the certainty of a measurement.

• Exact Numbers: Numbers known with complete certainty (e.g., counting numbers).

• Measured Numbers: Numbers obtained by measurement; contain uncertainty.

• Multiplication/Division: The result should have the same number of sig figs as the measurement with the fewest sig figs.

• Addition/Subtraction: The result should have the same number of decimal places as the measurement with the fewest decimal places.

Unit Conversions

Purpose and Method

Unit conversions are used to change measurements from one unit to another using conversion factors.

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit.

• Set-up: Arrange conversion factors so units cancel appropriately.

• Sequential Conversions: Multiple steps may be needed for complex conversions.

• Example: To convert 10 cm to meters: $10 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.1 \text{ m}$

Heat and Energy

Temperature and Heat

Temperature measures the average kinetic energy of particles, while heat is the energy transferred due to temperature difference.

• Temperature Units: Kelvin (K), Celsius (°C)

• Conversions: $T(\text{K}) = T(\text{°C}) + 273.15$

• Heat Units: Joule (J), calorie (cal)

• Specific Heat: The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Equation: $q = m \times c \times \Delta T$ where $q$ = heat, $m$ = mass, $c$ = specific heat, $\Delta T$ = change in temperature.

Density

Definitions and Applications

Density is a physical property that relates mass and volume.

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Specific Gravity: Ratio of the density of a substance to the density of water.

• Floating: Objects with lower density than the liquid will float.

• Hydrometer: Instrument used to measure specific gravity.

Chemical and Physical Changes

Types of Changes and Properties

Matter can undergo physical or chemical changes.

• Physical Changes: Changes in state or appearance without altering composition (e.g., melting, boiling).

• Physical Properties: Characteristics observed without changing the substance (e.g., color, density).

• Chemical Changes: Changes that produce new substances (e.g., rusting, burning).

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

States of Matter

Solid, Liquid, Gas

Matter exists in three primary states, each with distinct properties.

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

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

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

Classification of Matter

Pure Substances and Mixtures

Matter can be classified as pure substances or mixtures.

• Pure Substance: Has a fixed composition; can be an element or a compound.

• Element: Cannot be broken down into simpler substances.

• Compound: Composed of two or more elements chemically combined.

• Mixture: Combination of two or more substances not chemically combined.

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

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

Chemical Reactions (Chemical Changes)

Writing and Identifying Reactions

Chemical reactions involve the transformation of substances into new substances.

• Writing Reactions: Use chemical equations to represent reactants and products.

• Signs of Chemical Change: Color change, gas production, formation of a precipitate, energy change (heat/light).