Chapter 1: Units of Measurement, Physical and Chemical Change – A Molecular Approach

Chapter 1: Units of Measurement, Physical and Chemical Change

Introduction to Chemistry

Chemistry is a central science that seeks to understand the properties and behavior of matter by studying atoms and molecules. It connects with many other scientific disciplines and is fundamental to understanding the natural and material world.

• Chemistry: The science that investigates the composition, structure, properties, and changes of matter.

• Atoms: The smallest units of matter that retain the identity of an element.

• Molecules: Two or more atoms bonded together in specific geometric arrangements.

• Interdisciplinary Nature: Chemistry overlaps with physics, biology, earth sciences, materials science, and more.

Example: Chemistry is essential in fields such as medicine, environmental science, food science, and nanotechnology.

Classification of Sciences

Chemistry is part of the physical sciences and is interconnected with natural and human sciences. Many joint programs and interdisciplinary fields involve chemistry.

• Physical Sciences: Physics, Chemistry, Materials Science

• Natural Sciences: Biology, Earth Science, Geography, Marine Science

• Human Sciences: Anatomy, Neuroscience, Psychology, Pharmacology

Example: Computational chemistry combines chemistry with computer science and mathematics.

Atoms and Molecules

Atoms: The Building Blocks of Matter

Atoms are microscopic particles that serve as the fundamental building blocks of all matter. Each element is made up of one type of atom.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Chemical Alphabet: Just as the English alphabet forms words, the chemical alphabet (elements) forms compounds.

Molecules: Combinations of Atoms

Molecules are formed when two or more atoms are bonded together in a specific geometric arrangement. The properties of molecules depend on their composition and structure.

• Chemical Bonds: Forces that hold atoms together in molecules.

• Example: NaCl (table salt) is a compound formed from sodium and chlorine atoms.

• Different molecules can have different shapes, patterns, and properties (e.g., methanethiol vs. dimethyl sulfide).

Physical and Chemical Properties and Changes

Physical vs. Chemical Properties

Properties of matter can be classified as physical or chemical.

• Physical Property: Can be observed or measured without changing the substance's chemical identity (e.g., color, melting point, density).

• Chemical Property: Describes how a substance interacts with other substances, resulting in a change in chemical composition (e.g., flammability, reactivity with oxygen).

Example: The odor of hydrogen sulfide (rotten eggs) is a physical property; the ability of iron to rust is a chemical property.

Physical vs. Chemical Changes

Changes in matter are classified as physical or chemical.

• Physical Change: Alters the form or appearance of matter but does not change its chemical identity (e.g., melting, freezing, dissolving).

• Chemical Change: Involves one or more substances transforming into new substances with different chemical compositions (e.g., burning, rusting).

Law of Conservation of Mass: Mass is neither created nor destroyed during physical or chemical changes.

Signs of Chemical Change

• Bubbles of gas appear

• A precipitate (solid) forms

• Unexpected color change

• Change in melting or boiling point

Energy and Matter

Types of Energy

Physical and chemical changes are often accompanied by energy changes.

• Kinetic Energy (KE): Energy associated with the motion of an object.

• Formula: $KE = \frac{1}{2}mv^2$

• Thermal Energy: Associated with the temperature of an object, resulting from the motion of atoms and molecules.

• Potential Energy (PE): Energy due to position or composition.

• Formula: $PE = mgh$ (for gravitational potential energy)

• Chemical Energy: A form of potential energy stored in chemical bonds.

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

Measurement and Units

SI Units and Measurement

Quantitative observations in chemistry require standard units of measurement. The International System of Units (SI) is used in science.

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Amount of substance: mole (mol)

Mass vs. Weight: Mass is the amount of matter in an object; weight is the force of gravity on that mass.

Temperature Scales

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

• Kelvin (K): Absolute temperature scale; 0 K is absolute zero.

• Fahrenheit (°F): Water freezes at 32°F and boils at 212°F.

Conversion Formulas:

• $T_{K} = T_{°C} + 273.15$

• $T_{°F} = (T_{°C} \times \frac{9}{5}) + 32$

• $T_{°C} = (T_{°F} - 32) \times \frac{5}{9}$

Scientific Notation and Prefix Multipliers

Scientific notation is used to express very large or very small numbers conveniently. Prefix multipliers are used in the metric system to indicate powers of ten.

• Scientific Notation: $N \times 10^n$, where $1 \leq N < 10$ and $n$ is an integer.

• Example: $0.001 = 1 \times 10^{-3}$

• Common Prefixes: kilo- (k, $10^3$), centi- (c, $10^{-2}$), milli- (m, $10^{-3}$), micro- (μ, $10^{-6}$), nano- (n, $10^{-9}$)

Significant Figures

Significant figures (sig figs) indicate the precision of a measured quantity. The rules for determining significant figures are:

• 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 have an unlimited number of significant figures.

Rounding: Round to the correct number of significant figures based on the operation performed.

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

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

Problem Solving in Chemistry

Dimensional Analysis and Conversion Factors

Dimensional analysis uses conversion factors to convert from one unit to another. Conversion factors are ratios that express the relationship between two units.

• Example: To convert 15 miles to kilometers, use the conversion factor $1~\text{mile} = 1.609~\text{km}$.

• Set up the calculation so that units cancel appropriately.

Volume and Density

Volume is the amount of space occupied by a substance. Density is the ratio of mass to volume and is an intensive property (independent of amount).

• Volume (V): $V = \text{length}^3$

• SI unit: cubic meter ($m^3$); commonly used: liter (L), milliliter (mL), cubic centimeter ($cm^3$)

• Density (d): $d = \frac{m}{V}$

• Units: $g/cm^3$, $g/mL$ for solids and liquids; $g/L$ for gases

Example Table: Densities of Common Substances

Substance	Density (g/mL)
N-butyl alcohol	0.810
Ethylene glycol	1.114
Isopropyl alcohol	0.785
Toluene	0.866

Additional info: Density is temperature-dependent; heating generally decreases density.

Order of Operations

When performing calculations, follow the standard order of operations:

• Parentheses

• Exponents

• Multiplication and Division (left to right)

• Addition and Subtraction (left to right)

Problem-Solving Strategy

• Identify the given information.

• Determine what is to be found.

• Develop a conceptual plan (steps to follow).

• Carry out the plan, watching for significant figures and unit cancellation.

• Check if the answer makes sense.