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

Chapter 1: Units of Measurement for Physical and Chemical Change

Chapter Objectives

• Define, recognize, and distinguish between physical and chemical changes and properties.

• Understand the types of energy associated with these changes and the law of conservation of energy.

• Know and use the International System of Units (SI) for measurement, including SI prefixes and temperature scales.

• Determine significant figures in measurements and calculations; distinguish between accuracy and precision.

• Apply problem-solving strategies for unit conversion and algebraic manipulation in chemical problems.

Physical and Chemical Changes or Properties

Definition of Matter and Its Properties

Chemistry is the study of the properties, composition, and behavior of matter. Matter is anything that takes up space and has mass. The properties of matter are determined by the atoms or molecules that compose it.

Physical Properties

Physical property: A property that is displayed without changing the composition of the substance.

• Examples: Odour, taste, colour, appearance, melting point, boiling point, density.

Physical Changes

Physical change: A change in matter without a change in composition.

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

• Heating ice causes it to melt to liquid water and then evaporate to steam, but the molecular composition remains unchanged.

Chemical Properties

Chemical property: The ability of matter to undergo a change in composition.

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

Chemical Changes

Chemical change: A change in matter that results in a change in composition. Atoms rearrange to form new substances.

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

Energy Associated with Physical and Chemical Changes

Types of Energy

• Energy: The capacity to do work.

• Kinetic energy: Energy associated with motion.

• Potential energy: Energy associated with position or composition.

• Thermal energy: Energy associated with the temperature of an object (a form of kinetic energy).

Law of Conservation of Energy

Law of Conservation of Energy: Energy cannot be created or destroyed; it can only be transformed from one form to another.

This is the first law of thermodynamics:

$\text{Total Energy}_{\text{initial}} = \text{Total Energy}_{\text{final}}$

Energy in Physical and Chemical Changes

• Physical change example: Alcohol evaporating from skin absorbs heat, making the skin feel cooler (energy is absorbed).

• Chemical change example: Burning natural gas releases heat, which can be used for cooking or heating (energy is released).

The Units of Measurement

International System of Units (SI)

The SI system is a standardized set of units used in science:

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

SI Prefixes

SI prefixes are used to express multiples or fractions of units:

Prefix	Symbol	Multiplier	Power of Ten
tera	T	1,000,000,000,000	10^12
giga	G	1,000,000,000	10^9
mega	M	1,000,000	10^6
kilo	k	1,000	10^3
centi	c	0.01	10^-2
milli	m	0.001	10^-3
micro	μ	0.000001	10^-6
nano	n	0.000000001	10^-9

Temperature Scales

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

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

• Conversion:

  • $K = ^\circ C + 273.15$

  • $^\circ C = K - 273.15$

Derived Units

• Volume: $V = l \times w \times h$ (for a rectangular solid)

• 1 L = 1 dm3 = 1000 cm3

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

• Common units: g/mL, g/cm3, kg/m3

Special Derived Units

Quantity	Unit	Symbol	Definition
Energy	Joule	J	kg·m2/s2
Pressure	Pascal	Pa	kg/(m·s2)
Force	Newton	N	kg·m/s2
Charge	Coulomb	C	A·s
Voltage	Volt	V	J/C

The Reliability of a Measurement

Significant Figures

Significant figures reflect the precision of a measured quantity. The last digit is always estimated.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros after a decimal point are significant.

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

Scientific notation is used to clearly indicate significant figures (e.g., 1.20 × 103 has 3 significant figures).

Significant Figures in Calculations

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

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

• Logarithms: The number of significant figures in the result (after the decimal) matches 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 be random (affecting precision) or systematic (affecting accuracy).

Solving Chemical Problems

Problem-Solving Strategy

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

2. Strategize: Determine what you are asked to solve for.

3. Solve: Devise a plan to connect the given information to the unknown, which may involve unit conversions or algebraic manipulation.

4. Check: Ensure the answer makes sense and is reported with the correct units and significant figures.

Example: If given mass and volume, and asked for density, use $d = \frac{m}{V}$ and ensure units are consistent.