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

Introduction to Chemistry

The Central Science

Chemistry is often called the central science because it connects mathematics, physics, biology, earth science, and medicine. Understanding chemistry helps us comprehend the building blocks of matter and the reactions that power life, fuel technology, and shape the natural world. Applications range from developing life-saving drugs to creating sustainable materials.

• Chemistry provides tools to solve real-world problems.

• It is essential for innovations in medicine, energy, and materials science.

Matter and Energy

Definitions and Examples

The universe is composed of matter and energy. These two fundamental concepts are central to all chemical phenomena.

• Matter: Anything that has mass and takes up space (e.g., atoms, molecules, air, water, stars).

• Energy: The ability to do work (e.g., light, heat, motion, electricity).

Example: Water, sugar, and DNA are all forms of matter with unique molecular structures that determine their properties and functions.

Why Do Properties of Matter Matter?

Molecular Structure and Properties

The properties of matter are determined by the properties of molecules. The structure, shape, and interactions of molecules dictate how substances behave.

• Water: Bent shape and polarity make it a universal solvent with high heat capacity.

• Sugar: Structure allows it to be easily metabolized for energy.

• DNA: Double helix structure provides stable storage and transmission of genetic information.

Physical and Chemical Properties of Matter

Definitions and Applications

Chemistry seeks to understand the properties and behavior of matter by studying atoms and molecules. Recognizing these properties helps us:

• Identify substances: Each material has a unique fingerprint of physical and chemical traits.

• Predict behavior: Will it burn, dissolve, react with acid, or remain stable?

• Choose the right material: For building, cooking, medicine, or technology.

Physical and Chemical Changes

Distinguishing Changes

Physical changes alter the appearance or state of a substance without forming new substances. Chemical changes result in the formation of new substances through chemical reactions.

• Physical Change: No new substances; only the appearance or state changes.

• Chemical Change: New substances are formed; a chemical reaction occurs.

Examples:

• Sugar dissolving in water – Physical change (molecules remain unchanged).

• Wood burning – Chemical change (new substances like ash and carbon dioxide are formed).

• Water freezing – Physical change (state changes, but still H2O).

• Egg cooking – Chemical change (proteins change structure).

• Soda losing its fizz – Physical change (CO2 escapes, soda remains).

Energy in Physical and Chemical Change

Law of Conservation of Energy

Energy is a fundamental part of all changes in chemistry. The law of conservation of energy states that energy can change forms but cannot be created or destroyed.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy.

Equation:

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

Example: In a car engine, chemical energy in gasoline is converted to thermal energy (heat) and then to mechanical energy (motion).

Units of Measurement

SI Base Units

Scientific measurements use the International System of Units (SI), which includes seven base units:

Quantity	Unit Name	Symbol
Length	metre	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Electric current	ampere	A
Luminous intensity	candela	cd

Example: The CN Tower is 553.3 metres tall; a hockey player is about 2 metres tall.

Mass and Weight

• Mass: Quantity of matter in an object (measured in kilograms or grams).

• Weight: Gravitational force acting on an object due to its mass.

Time

• Second: Defined by the duration of 9,192,631,770 periods of radiation from a cesium-133 atom.

Temperature

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

• Celsius (°C): Commonly used in daily life. Conversion: $T(K) = T(°C) + 273.15$

Derived Units

Derived units are combinations of base units, commonly used in chemistry.

Quantity	Unit Name	Symbol
Volume	cubic metre	m3
Density	kilogram per cubic metre	kg/m3
Speed	metre per second	m/s

Example: 1 litre (L) = 1 dm3 = 1000 mL = 1000 cm3

Metric Prefixes

Prefixes are used to express multiples or fractions of units.

Prefix	Symbol	Multiplier
giga	G	109
mega	M	106
kilo	k	103
centi	c	10-2
milli	m	10-3
micro	μ	10-6
nano	n	10-9

Example: $1.25 \times 10^{-1} \text{ L} = 125 \text{ mL}$

Reporting Measurements

Significant Figures

Significant figures reflect the precision of a measurement. All nonzero digits are significant; zeros may or may not be, depending on their position.

• Interior zeros (between nonzero digits) are significant.

• Leading zeros (before nonzero digits) are not significant.

• Trailing zeros after a decimal point are significant.

• Trailing zeros before an implied decimal point are ambiguous; use scientific notation.

Example: 45.000 (5 sig figs), 0.0032 (2 sig figs), 1.200 × 102 (4 sig figs)

Significant Figures in Calculations

• Multiplication/Division: Result has the same number of significant figures as the factor with the fewest significant figures.

• Addition/Subtraction: Result has the same number of decimal places as the quantity with the fewest decimal places.

• Logarithms: The mantissa (digits after the decimal) in the result matches the number of significant figures in the original number.

• Antilogarithms: The number of significant figures in the result equals the number of digits in the mantissa.

Precision and Accuracy

Definitions

• Accuracy: How close a measurement is to the true or accepted value (hitting the bullseye).

• Precision: How consistent repeated measurements are (hitting the same spot repeatedly).

Example: Three students weigh a block known to have a mass of 10.00 g. Their results may be precise (close to each other) or accurate (close to the true value).

Intensive vs. Extensive Properties

Classification of Properties

Property Type	Definition	Depends on Amount?	Examples
Intensive	Does not depend on the size or amount of substance	No	Temperature, boiling point, color, conductivity, density
Extensive	Depends on the size or amount of substance	Yes	Mass, volume, length, energy, heat capacity

Density

Definition and Calculation

Density is the mass per unit volume of a substance. It indicates how much matter is packed into a given space.

Equation:

$d = \frac{m}{V}$

• m: mass (g or kg)

• V: volume (mL, cm3, or m3)

Example: A metal sample has a mass of 87.5 g and displaces water from 45.0 mL to 55.8 mL. Volume = 55.8 mL - 45.0 mL = 10.8 mL. Density = $\frac{87.5\,\text{g}}{10.8\,\text{mL}} = 8.10\,\text{g/mL}$

To convert to kg/m3: $8.10\,\text{g/mL} \times 1000 = 8100\,\text{kg/m}^3$

Measurement Uncertainty

Precision in Lab Measurements

All measurements have uncertainty. Instruments are not perfectly exact; always report one estimated digit beyond what is certain.

• Analog instruments: Estimate between marked values.

• Digital instruments: Last digit is the estimated digit.

Example: A ruler reading of 54.7 mm ± 0.1 mm; a graduated cylinder reading of 23.6 mL ± 0.1 mL.

Solving Chemical Problems

Unit Conversions and Reporting

Unit conversions are essential for reporting measurements in science. Prefixes and scientific notation help communicate values clearly and avoid errors.

• Convert between units using conversion factors.

• Report measurements using appropriate significant figures and units.

Example: $0.0015\,\text{L} = 1.5\,\text{mL}$; $0.00000000085\,\text{m} = 0.85\,\text{nm}$

Additional info: Some tables and examples have been expanded for clarity and completeness.