Chapter 1: Matter and Measurements – GOB Chemistry Study Notes

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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—anything that has mass and occupies space—and the changes it undergoes.

Chemistry: The study of the nature, properties, and transformations of matter.
Matter: Physical material that makes up the universe; has mass and occupies space.
Scientific Method: A systematic process involving observation, hypothesis, and experimentation to expand scientific knowledge.
Property: A characteristic useful for identifying a substance (e.g., size, color, temperature, chemical composition, chemical reactivity).
Physical Change: Alters the form or appearance of matter but does not change its composition (e.g., melting, dissolving).
Chemical Change: Alters the chemical composition, resulting in 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, each with distinct properties.

Solid (s): Definite shape and volume.
Liquid (l): Definite volume, takes the shape of its container.
Gas (g): No definite shape or volume; expands to fill its container.
Change of State: Conversion from one state to another (e.g., melting, boiling).

Example: At room temperature, formaldehyde is a gas because its boiling point is below room temperature.

Classification of Matter

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

Pure Substance: Uniform chemical composition throughout; can be an element or a compound.
Mixture: Blend of two or more substances, each retaining its own identity.
Homogeneous Mixture: Uniform composition throughout (e.g., salt water).
Heterogeneous Mixture: Non-uniform composition (e.g., salad, granite).
Element: Fundamental substance that cannot be broken down chemically.
Compound: Pure substance that can be broken down into simpler substances by chemical reactions.

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

Chemical Elements and Symbols

Elements are represented by symbols, and their combinations form chemical formulas.

There are 118 known elements; 91 occur naturally.
Element Symbols: One or two letters; first letter capitalized, second (if any) lowercase (e.g., H for hydrogen, Na for sodium).
Some symbols are based on Latin names (e.g., Na for natrium/sodium).
Chemical Formula: Notation using element symbols and subscripts to indicate the number of atoms (e.g., H2O).
Elements essential for human life include C, H, O, N, and others, classified as macronutrients or micronutrients based on required amounts.

Element	Symbol	Latin Name (if applicable)
Sodium	Na	Natrium
Potassium	K	Kalium
Iron	Fe	Ferrum
Copper	Cu	Cuprum
Lead	Pb	Plumbum

Additional info: Table includes only a few examples; refer to periodic table for full list.

Chemical Reactions: Examples of Chemical Change

Chemical reactions are processes in which substances (reactants) are transformed into new substances (products).

Chemical Reaction: Process changing the identity and composition of substances.
Reactant: Starting material in a chemical reaction.
Product: Substance formed as a result of a chemical reaction.
Reactions are represented with reactants on the left, products on the right, and an arrow indicating the direction of change.

Example: Electrolysis of water produces hydrogen and oxygen gases.

Physical Quantities: Units and Scientific Notation

Physical properties are measured using standardized units, often expressed in scientific notation for convenience.

Physical Quantity: Measurable property (e.g., mass, volume, temperature).
Unit: Standard of measurement (e.g., kilogram, meter, liter, kelvin, second).
SI Units: International System of Units; base units include kg (mass), m (length), K (temperature), s (time), m3 (volume).
Metric Units: Gram (g), liter (L), Celsius (°C).
Derived Units: Formed from base units (e.g., speed in m/s, density in g/cm3).
Prefixes: Used to indicate multiples or fractions of units (e.g., milli-, micro-, kilo-).
Scientific Notation: Expresses numbers as a product of a number between 1 and 10 and a power of 10.

Prefix	Symbol	Multiplier
Kilo	k	103
Milli	m	10-3
Micro	μ	10-6
Nano	n	10-9

Example: 0.000000120 m = m = 120 nm

Measuring Mass, Length, and Volume

Standard units are used to measure mass, length, and volume in chemistry.

Mass: Amount of matter; measured in kilograms (kg) or grams (g).
Weight: Gravitational force on an object; differs from mass.
Length: Measured in meters (m), centimeters (cm), or millimeters (mm).
Volume: Space occupied; measured in liters (L), milliliters (mL), or cubic centimeters (cm3).

Quantity	SI Unit	Common Metric Unit
Mass	kg	g
Length	m	cm, mm
Volume	m3	L, mL, cm3

Measurement and Significant Figures

All measurements have some uncertainty, reflected in the number of significant figures reported.

Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.
Rules:
- Zeroes between nonzero digits are significant.
- Leading zeroes are not significant.
- Trailing zeroes after a decimal point are significant.
- Trailing zeroes before an implied decimal point may or may not be significant.
- Exact numbers (e.g., 1 foot = 12 inches) have infinite significant figures.

Example: 0.0076 mL has two significant figures; 3400 kg may have two, three, or four significant figures depending on context.

Rounding Off Numbers

Rounding is used to limit the number of significant figures in a result to reflect measurement precision.

Multiplication/Division: Result has as many significant figures as the least precise measurement.
Addition/Subtraction: Result has as many decimal places as the least precise measurement.
Rounding Rules:
- If the first digit dropped is 4 or less, drop it and all following digits.
- If the first digit dropped is 5 or greater, increase the last retained digit by 1.

Example: 4.5832 rounded to two significant figures is 4.6.

Problem Solving: Unit Conversions and Estimating Answers

Unit conversions are essential in chemistry and are often performed using the factor-label (dimensional analysis) method.

Conversion Factor: Ratio expressing the relationship between two units; always equals 1.
Factor-Label Method: Set up calculations so that unwanted units cancel, leaving the desired unit.
Steps:
1. Identify given information and units.
2. Identify the desired answer and units.
3. Find relationships (conversion factors) between known and unknown units.
4. Solve, checking that units cancel appropriately.
Ballpark Estimate: Make a rough estimate to check if the answer is reasonable.

Example: To convert 21.5 inches to centimeters, use the conversion factor 1 in = 2.54 cm.

Temperature, Heat, and Energy

Temperature and energy are fundamental concepts in chemistry, with several scales and units in use.

Energy: Capacity to do work or supply heat; measured in joules (J) or calories (cal).
Temperature: Measure of heat energy; SI unit is kelvin (K), but Celsius (°C) and Fahrenheit (°F) are also used.
Temperature Conversions:
Calorie: Amount of heat needed to raise 1 g of water by 1°C.
Specific Heat: Amount of heat required to raise the temperature of 1 g of a substance by 1°C.
Heat Calculation: where = heat (cal or J), = mass (g), = specific heat, = temperature change (°C)

Example: Heating 95 kg of water from 18°C to 43°C requires cal.

Density and Specific Gravity

Density relates mass and volume, and specific gravity compares the density of a substance to that of water.

Density: Mass per unit volume;
Units: g/cm3 for solids, g/mL for liquids.
Specific Gravity: Ratio of the density of a substance to the density of water at the same temperature.
Substances with density less than water float; those with greater density sink.
Measured using a hydrometer; in medicine, a urinometer is used for urine.

Example: To find the volume of 25.0 g of isopropyl alcohol (density = 0.7855 g/mL):

Summary Table: Key Physical Quantities and Units

Quantity	SI Unit	Common Metric Unit	Conversion
Mass	kg	g	1 kg = 1000 g
Length	m	cm, mm	1 m = 100 cm = 1000 mm
Volume	m3	L, mL	1 L = 1000 mL = 1000 cm3
Temperature	K	°C	K = °C + 273.15

Concept Map

Matter → Pure Substance (Element or Compound) or Mixture (Homogeneous or Heterogeneous) → Physical and Chemical Properties → Physical and Chemical Changes