Chapter 1: Matter and Measurements – GOB Chemistry Study Notes

Matter and Measurements

Chemistry: The Central Science

Chemistry is known as the central science because it connects and underpins all other scientific disciplines. It focuses on the study of matter, its properties, and the changes it undergoes.

• Chemistry: The study of the nature, properties, and transformations of matter.

• Matter: Anything that has mass and occupies space.

• Scientific Method: A systematic process involving observation, hypothesis, and experimentation.

• Property: A characteristic used to identify substances (e.g., size, color, temperature, chemical composition, chemical reactivity).

• Physical Change: Alters the form or appearance but not the chemical identity (e.g., melting, dissolving).

• Chemical Change: Alters the chemical identity, forming new substances (e.g., burning, rusting).

Example: Dissolving sugar in water is a physical change; heating sugar to make caramel is a chemical change.

States of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct physical properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes the shape of its container.

• Gas: No definite shape or volume; fills the container.

• Change of State: Transition between states (e.g., melting, boiling).

Example: Formaldehyde is a gas at room temperature because room temperature is above its boiling point.

Classification of Matter

All matter can be classified as either a pure substance or a mixture.

• Pure Substance: Uniform chemical composition throughout.

• Mixture: Blend of two or more substances, each retaining its identity.

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

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

• Element: Cannot be broken down chemically; fundamental substance.

• Compound: Can be broken down chemically; composed of two or more elements.

Example: Vanilla ice cream is a homogeneous mixture; sugar is a compound.

Chemical Elements and Symbols

Elements are represented by one- or two-letter symbols, often derived from their English or Latin names.

• Element Symbols: First letter capitalized, second letter lowercase (e.g., Na for sodium).

• Chemical Formula: Notation showing the elements and their ratios in a compound (e.g., H2O).

• Subscripts: Indicate the number of atoms of each element.

• Macronutrients: Essential elements required in amounts >100 mg/day.

• Micronutrients: Essential elements required in amounts ≤15 mg/day.

Element	Symbol	Latin Name
Sodium	Na	Natrium
Potassium	K	Kalium
Iron	Fe	Ferrum
Gold	Au	Aurum
Silver	Ag	Argentum
Copper	Cu	Cuprum
Lead	Pb	Plumbum
Tin	Sn	Stannum

Additional info: Table inferred from common element symbols and Latin names.

Chemical Reactions: Examples of Chemical Change

Chemical reactions are represented using symbolic notation, showing reactants and products.

• Chemical Reaction: Process where substances change identity and composition.

• Reactant: Starting substance.

• Product: Substance formed.

• Notation: Reactants on the left, products on the right, arrow indicates reaction.

Example: Electrolysis of water produces hydrogen and oxygen.

$\mathrm{2H_2O(l) \xrightarrow{electricity} 2H_2(g) + O_2(g)}$

Physical Quantities: Units and Scientific Notation

Physical properties are measured using units. Scientific notation is used for very large or small numbers.

• Physical Quantity: Measurable property (e.g., mass, volume).

• Unit: Standard of measurement (e.g., kg, m, L, K, s).

• SI Units: International System of Units.

• Metric Prefixes: Modify unit size (e.g., milli-, micro-, kilo-).

• Scientific Notation: $a \times 10^n$ where $1 \leq a < 10$.

Prefix	Symbol	Multiplier
Kilo	k	$10^3$
Milli	m	$10^{-3}$
Micro	μ	$10^{-6}$
Nano	n	$10^{-9}$

Example: 0.000000120 m = $1.2 \times 10^{-7}$ m = 120 nm

Measuring Mass, Length, and Volume

Mass, length, and volume are fundamental measurements in chemistry.

• Mass: Amount of matter (kg, g).

• Weight: Gravitational force on an object.

• Length: Distance (m, cm).

• Volume: Space occupied (L, mL, cm3).

Quantity	SI Unit	Metric Unit
Mass	kg	g
Length	m	cm
Volume	m3	L, mL

Example: 1 L = 1000 mL; 1 m = 100 cm

Measurement and Significant Figures

Significant figures indicate the precision of a measurement. Rules determine which digits are significant.

• Significant Figures: All certain digits plus one estimated digit.

• Rule 1: Zeroes in the middle are significant (e.g., 94.072 g).

• Rule 2: Leading zeroes are not significant (e.g., 0.0834 cm).

• Rule 3: Trailing zeroes after a decimal are significant (e.g., 138.200 m).

• Rule 4: Trailing zeroes before an implied decimal may or may not be significant (e.g., 23,000 kg).

• Exact Numbers: Have unlimited significant figures (e.g., 1 foot = 12 inches).

Example: 2730.78 m (6 sig figs), 0.0076 mL (2 sig figs), 3400 kg (2-4 sig figs), 3400.0 (5 sig figs)

Rounding Off Numbers

Rounding is used to limit the number of significant figures in calculated results.

• Multiplication/Division: Result has no more sig figs than the least precise number.

• Addition/Subtraction: Result has no more decimal places than the least precise number.

• Rule 1: If the first dropped digit is 4 or less, drop it.

• Rule 2: If the first dropped digit is 5 or greater, round up.

Example: 2.4271 rounded to two sig figs is 2.4; 4.5832 rounded to two sig figs is 4.6.

Problem Solving: Unit Conversions and Estimating Answers

The factor-label method (dimensional analysis) is used to convert between units.

• Conversion Factor: Numerical relationship between two units (e.g., 1 lb = 454 g).

• Steps:

  1. Identify given information and units.

  2. Identify desired answer and units.

  3. Find relationships and plan steps.

  4. Solve and check with a ballpark estimate.

Example: To convert 21.5 inches to centimeters: $21.5 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 54.6 \text{ cm}$

Temperature, Heat, and Energy

Temperature measures heat energy. Different scales are used: Celsius, Kelvin, and Fahrenheit.

• Energy: Capacity to do work or supply heat (SI unit: joule, J; metric unit: calorie, cal).

• Temperature: Measured in K (SI), °C (metric), or °F (imperial).

• Celsius: Freezing point of water is 0°C, boiling point is 100°C.

• Kelvin: Absolute zero is 0 K; same degree size as Celsius.

• Fahrenheit: Freezing point is 32°F, boiling point is 212°F.

• Conversion Equations:

$\mathrm{T_{C} = \frac{5}{9}(T_{F} - 32)}$

$\mathrm{T_{K} = T_{C} + 273.15}$

• Specific Heat: Amount of heat needed to raise 1 g of a substance by 1°C.

• Heat Calculation:

$\mathrm{q = m \times c \times \Delta T}$

where $q$ = heat (cal or J), $m$ = mass (g), $c$ = specific heat (cal/g°C), $\Delta T$ = temperature change (°C).

Example: Heating 95 kg of water by 25°C requires $q = 95,000 \text{ g} \times 1 \text{ cal/g°C} \times 25 \text{°C} = 2,375,000 \text{ cal}$

Density and Specific Gravity

Density relates mass to volume; specific gravity compares density to water.

• Density: $\mathrm{Density = \frac{mass}{volume}}$

• Units: g/cm3 (solids), g/mL (liquids).

• Specific Gravity: $\mathrm{Specific\ Gravity = \frac{density\ of\ substance}{density\ of\ water}}$

• Hydrometer: Instrument to measure specific gravity.

• Floating/Sinking: Less dense substances float; more dense substances sink.

Example: To find the volume of 25.0 g of isopropyl alcohol (density = 0.7855 g/mL): $\mathrm{Volume = \frac{mass}{density} = \frac{25.0\ g}{0.7855\ g/mL} = 31.8\ mL}$

Material	Density (g/cm3)
Water	1.00
Isopropyl Alcohol	0.7855
Ice	0.92
Iron	7.87
Gold	19.3

Additional info: Table inferred from common densities.

Concept Map

A concept map visually organizes the relationships between matter, measurement, and their properties, connecting topics such as classification, states, elements, reactions, units, and calculations.