General Chemistry I - CHM 101: Chapter 1 Study Notes – Matter, Energy & Measurement

Chapter 1: Matter, Energy & Measurement

Overview

This chapter introduces the foundational concepts of chemistry, focusing on the classification of matter, its properties, the nature of energy, and the principles of measurement and unit conversion. Mastery of these topics is essential for understanding all subsequent material in general chemistry.

Matter: Classification and Composition

Definition of Matter

• Matter is anything that has mass and occupies space.

Classification by Composition

• Substance: A form of matter with distinct properties and a composition that does not vary from sample to sample.

• Mixture: A combination of two or more substances in which each retains its own chemical identity and properties.

Types of Substances

• Element: A substance that cannot be decomposed into simpler substances. Elements are made up of identical atoms, which are the fundamental building blocks of matter.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions. Compounds have molecules containing atoms of different elements.

Law of Constant Composition (Law of Definite Proportions)

• This law states that the elemental composition of a compound is always the same, regardless of its source or how it was prepared.

• Example: One molecule of methane (CH4) always contains one carbon atom and four hydrogen atoms.

Types of Mixtures

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

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

States of Matter

Overview of States

• Matter exists in three primary states: gas, liquid, and solid, each with distinct observable properties.

Properties of Each State

• Gas: No fixed shape or volume; fills the container; particles are far apart and move rapidly; easily compressed or expanded.

• Liquid: Fixed volume but no fixed shape; takes the shape of the container; particles are close together but can move past one another; not easily compressed.

• Solid: Definite shape and volume; particles are closely packed in fixed positions; not compressible.

Properties and Changes of Matter

Chemical vs. Physical Properties

• Physical Properties: Can be observed without changing the substance into another substance (e.g., boiling point, density, mass, volume, odor, hardness).

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

Chemical vs. Physical Changes

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

• Chemical Change: Changes that result in the formation of new substances (e.g., combustion, oxidation, decomposition).

• Example: Melting ice is a physical change; burning hydrogen in oxygen to form water is a chemical change.

Intensive vs. Extensive Properties

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

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

Energy in Chemistry

Types of Energy

• Kinetic Energy: The energy of motion. The faster an object moves, the greater its kinetic energy.

• Potential Energy: Stored energy due to an object's position or arrangement.

Formula for Kinetic Energy:

$KE = \frac{1}{2}mv^2$

• Where m is mass and v is velocity.

Measurement and Units

SI Base Units

• The International System of Units (SI) is the standard for scientific measurements.

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

Metric Prefixes

Metric prefixes are used to indicate multiples or fractions of units.

Prefix	Abbreviation	Factor
Peta	P	$10^{15}$
Tera	T	$10^{12}$
Giga	G	$10^{9}$
Mega	M	$10^{6}$
Kilo	k	$10^{3}$
Deci	d	$10^{-1}$
Centi	c	$10^{-2}$
Milli	m	$10^{-3}$
Micro	μ	$10^{-6}$
Nano	n	$10^{-9}$
Pico	p	$10^{-12}$
Femto	f	$10^{-15}$
Atto	a	$10^{-18}$
Zepto	z	$10^{-21}$

Temperature Scales

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

• Kelvin (K): The SI unit of temperature; absolute zero (0 K) is the lowest possible temperature.

• Fahrenheit (°F): Commonly used in the United States.

Conversion Formulas:

• $K = ^\circ C + 273.15$

• $^\circ F = \frac{9}{5}(^\circ C) + 32$

Volume

• Volume is a derived unit (length cubed). Common units: liter (L), milliliter (mL), cubic centimeter (cm3).

• $1\,\text{L} = 1\,\text{dm}^3 = 1000\,\text{mL} = 1000\,\text{cm}^3$

Energy Units

• Joule (J): The SI unit of energy. $1\,\text{J} = 1\,\text{kg} \cdot \text{m}^2/\text{s}^2$

• Calorie (cal): Commonly used in chemistry. $1\,\text{cal} = 4.184\,\text{J}$

Scientific Notation

Purpose and Structure

• Scientific notation expresses numbers as a product of a number between 1 and 10 and a power of 10.

• Example: $3.0 \times 10^4$

Rules for Scientific Notation

• If the original number is less than 1, the exponent is negative.

• If the original number is greater than 1, the exponent is positive.

Operations with Scientific Notation

• Multiplication: Multiply the base numbers and add the exponents.

• Division: Divide the base numbers and subtract the exponents.

• Addition/Subtraction: Convert to the same exponent before adding or subtracting the base numbers.

Significant Figures

Definition and Importance

• Significant figures (sig figs) are the digits in a measurement that are known with certainty plus one digit that is estimated.

• They reflect the precision of a measurement and prevent overstating accuracy in calculations.

Rules for Counting Significant Figures

• Non-zero digits are always significant.

• Leading zeros are not significant.

• Captive zeros (between non-zero digits) are significant.

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

• Exact numbers (from counting or definitions) have infinite significant figures.

Significant Figures in Calculations

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

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

Dimensional Analysis (Unit Conversion)

Principles

• Dimensional analysis uses conversion factors to convert from one unit to another.

• Set up the calculation so that units cancel, leaving only the desired unit.

Example: Single-Step Conversion

• Convert 3.000 m to cm:

$3.000\,\text{m} \times \frac{100\,\text{cm}}{1\,\text{m}} = 300.0\,\text{cm}$

Example: Multi-Step Conversion

• Convert 3.000 m to inches:

$3.000\,\text{m} \times \frac{100\,\text{cm}}{1\,\text{m}} \times \frac{1\,\text{in}}{2.54\,\text{cm}} = 118.1\,\text{in}$

Common Conversion Factors

Quantity	Conversion
Length	1 in = 2.54 cm (exact); 1 mi = 1.6093 km
Mass	1 kg = 2.2046 lb
Volume	1 L = 1000 cm3; 1 gal = 3.7854 L

Accuracy and Precision

Definitions

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

• Precision: How close repeated measurements are to each other.

Types of Numbers in Measurements

• Exact Numbers: Counted or defined values (e.g., 12 eggs in a dozen).

• Inexact Numbers: Measured values, subject to uncertainty.

Uncertainty in Measurements

• All measurements have some degree of uncertainty, depending on the instrument used.

Summary Table: Classification of Matter

Type	Definition	Example
Element	Cannot be decomposed into simpler substances	O2, Fe
Compound	Composed of two or more elements in fixed proportions	H2O, CO2
Homogeneous Mixture	Uniform composition throughout	Salt water, air
Heterogeneous Mixture	Non-uniform composition	Salad, potting soil

Additional info: Some context and examples have been expanded for clarity and completeness, as is standard in academic study guides.