Chapter 1: Units of Measurement for Physical and Chemical Change – Study Notes

Physical and Chemical Changes and Properties

Introduction to Matter and Its Properties

Chemistry is the study of the properties, composition, and behavior of matter. Matter is anything that occupies space and has mass. Understanding the properties of atoms and molecules that make up matter is fundamental to chemistry, as changes in matter arise from their physical and chemical properties.

• Physical Property: A characteristic of a substance that can be observed or measured without changing its chemical composition.

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

Physical Changes

A physical change is a change in matter that does not alter its chemical composition. The identity of the substance remains the same, even if its form or appearance changes.

• Example: Phase transitions of water (solid, liquid, gas). The composition of water is always H2O, regardless of its state.

• Heating water causes it to change from ice (solid) to liquid water and then to steam (gas), but the molecular composition remains unchanged.

Chemical Properties and Changes

A chemical property is the ability of a substance to undergo a change that alters its chemical composition.

• Examples: Combustibility of gasoline (reacts with oxygen to form new products), electrochemical nature of iron (can undergo oxidation to form rust).

A chemical change (chemical reaction) results in the formation of one or more new substances with different compositions and properties.

• Example: Iron reacts with oxygen to form iron oxide (rust).

Energy Associated with Physical and Chemical Changes

Energy in Chemical and Physical Processes

All physical and chemical changes are accompanied by changes in energy. Energy is the capacity to do work or transfer heat.

• Energy associated with a physical change: When alcohol evaporates from your skin, it absorbs heat from your body, making your skin feel cooler.

• Energy associated with a chemical change: Burning natural gas releases heat energy that can be used for cooking or heating.

Types of Energy

• Kinetic Energy: The energy of motion. For example, a falling object converts potential energy to kinetic energy.

• Potential Energy: The energy stored due to an object's position or composition. For example, a weight held at a height or the arrangement of atoms in a molecule.

Law of Conservation of Energy

The law of conservation of energy states that energy cannot be created or destroyed; it can only be transformed from one form to another. This is the first law of thermodynamics.

• In chemical reactions, energy is conserved but may be transferred as heat or work.

The Units of Measurement

International System of Units (SI)

The SI system is a standardized set of units used in science for consistency and clarity. The main SI base units relevant to chemistry are:

Quantity	Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K

Temperature Scales

• Celsius (°C): Commonly used in daily life and laboratory settings.

• Kelvin (K): The SI unit for temperature. Absolute zero (0 K) is the lowest possible temperature.

• Conversion: $T_{K} = T_{°C} + 273.15$

SI Prefixes

SI prefixes are used to express multiples or fractions of units, making it easier to handle very large or small numbers.

Prefix	Symbol	Multiplier	Power of Ten
giga	G	1,000,000,000	109
mega	M	1,000,000	106
kilo	k	1,000	103
centi	c	0.01	10-2
milli	m	0.001	10-3
micro	μ	0.000001	10-6
nano	n	0.000000001	10-9

Derived Units

• Volume: The amount of space occupied by a substance. Common units: liter (L), cubic meter (m3), cubic centimeter (cm3 or cc).

• Relationship: $1\,L = 1000\,cm^3$

• Density: Mass per unit volume. $\text{density} = \frac{\text{mass}}{\text{volume}}$

• Common units: g/mL, g/cm3, kg/m3

Substance	Density (g/mL or g/cm3)
Water	1.00
Ethanol	0.789
Copper	8.96
Gold	19.3
Lead	11.4
Mercury	13.55

• Concentration: Amount of solute per unit volume of solution. Commonly expressed as molarity (M): $c = \frac{n}{V}$, where $n$ is moles and $V$ is volume in liters.

Special Derived Units

• Pressure: Force per unit area. $P = \frac{F}{A}$

• SI unit: Pascal (Pa), where $1\,Pa = 1\,N/m^2$

• Other units: bar ($1\,bar = 10^5\,Pa$), atmosphere ($1\,atm = 101325\,Pa$), torr ($1\,torr = 133.322\,Pa$)

• Energy: Joule (J), $1\,J = 1\,kg\,m^2/s^2$

• Force: Newton (N), $1\,N = 1\,kg\,m/s^2$

The Reliability of a Measurement

Significant Figures

Significant figures (sig figs) indicate the precision of a measured quantity. The last digit in a measurement is always estimated.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

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

• Exact numbers (e.g., 100 cm in 1 m) have infinite significant figures.

Rules for Calculations

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

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

• Scientific Notation: Used to clearly indicate the number of significant figures (e.g., $1.20 \times 10^3$ has three significant figures).

• Logarithms: The number of significant figures in the result after the decimal point should match the number of significant figures in the original number.

Accuracy and Precision

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

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

• Errors can arise from poor technique (systematic error) or random fluctuations (random error).

Solving Chemical Problems

Problem-Solving Strategy

Effective problem-solving in chemistry involves a systematic approach:

1. Sort: Identify the information given and write down known values, constants, and relevant equations.

2. Strategize: Determine what you are asked to solve for and plan the steps needed to reach the solution.

3. Solve: Carry out the calculations, ensuring correct use of units and significant figures.

4. Check: Review your answer for reasonableness and correct units.

• Unit conversions and algebraic manipulation of equations are common in chemical problem-solving.

Example: Calculating Density

• Given: Mass and volume of a substance.

• To Find: Density using $\text{density} = \frac{\text{mass}}{\text{volume}}$

• Process: Substitute values, perform calculation, and report answer with correct significant figures.