Chapter 1: Matter & Measurements – GOB Chemistry Study Notes

Chapter 1: Matter & Measurements

Introduction to Chemistry

Chemistry is the scientific study of the composition, structure, properties, and reactions of matter. Understanding chemistry involves recognizing the different forms and classifications of matter, as well as the methods used to measure and describe it.

• Chemistry: The study of the composition, structure, properties, and reactions of matter.

• Matter: Anything that has mass and occupies space.

States of Matter

Matter exists in three primary physical states: solid, liquid, and gas. Each state has distinct characteristics regarding shape and volume.

• Solid: Has a definite shape and volume. Particles are closely packed in a fixed arrangement.

• Liquid: Has a definite volume but takes the shape of its container. Particles are close but can move past one another.

• Gas: Has neither definite shape nor volume. Particles are far apart and move freely.

Changes of State: Matter can change from one state to another through physical processes such as melting, freezing, vaporization, condensation, sublimation, and deposition.

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Physical and Chemical Properties

Properties of matter can be classified as physical or chemical.

• Physical Properties: Characteristics that can be observed or measured without changing the substance's identity (e.g., color, melting point, density).

• Chemical Properties: Characteristics that describe a substance's ability to undergo a change that transforms it into a different substance (e.g., flammability, reactivity).

Physical and Chemical Changes

Changes in matter are categorized as physical or chemical.

• Physical Change: A change that does not alter the identity of a substance (e.g., freezing, melting, dissolving).

• Chemical Change: A change that results in the formation of one or more new substances (e.g., burning, rusting).

Classification of Matter

Matter can be classified based on its composition as pure substances or mixtures.

• Pure Substance: Has a fixed composition and distinct properties. Can be an element or a compound.

• Element: A substance that cannot be broken down into simpler substances by chemical means (e.g., oxygen, gold).

• Compound: A substance composed of two or more elements chemically combined in fixed proportions (e.g., water, H2O).

• Mixture: A combination of two or more substances in which each retains its own properties. Can be homogeneous or heterogeneous.

Type	Description	Examples
Homogeneous Mixture	Uniform composition throughout; also called a solution	Salt water, air
Heterogeneous Mixture	Non-uniform composition	Oil & vinegar, pizza

Chemical Elements and Symbols

Elements are represented by unique symbols, usually one or two letters, with the first letter capitalized. Some symbols are derived from Latin names.

• Examples: H (Hydrogen), C (Carbon), Si (Silicon), Na (Sodium), Fe (Iron), Au (Gold)

Chemical Formulas

Compounds are represented by chemical formulas, which use element symbols and subscripts to indicate the number of atoms of each element present.

• Example: H2O (water) contains 2 hydrogen atoms and 1 oxygen atom.

• Example: C12H22O11 (sugar) contains 12 carbon, 22 hydrogen, and 11 oxygen atoms.

Measurements and Units

Physical quantities in chemistry are measured using standardized units. The main systems are the International System of Units (SI), the Metric System, and the United States Customary System (USCS).

• Mass: SI unit is kilogram (kg); metric unit is gram (g)

• Length: SI and metric unit is meter (m)

• Volume: SI unit is cubic meter (m3); metric unit is liter (L)

• Time: Second (s)

• Temperature: Kelvin (K), Celsius (°C)

• Energy: Joule (J), calorie (cal)

Common Unit Equivalents

Quantity	SI Unit	Metric Unit	Equivalents
Mass	kg	g	1 kg = 1000 g; 1 kg = 2.205 lb
Length	m	m	1 m = 100 cm; 1 in = 2.54 cm
Volume	m3	L	1 L = 1000 mL; 1 mL = 1 cm3

Prefixes for SI and Metric Units

Prefixes are used to indicate multiples or fractions of base units.

Prefix	Symbol	Multiplier	Example
kilo	k	103	1 kilometer (km) = 1000 meters (m)
centi	c	10-2	1 centimeter (cm) = 0.01 meter (m)
milli	m	10-3	1 milligram (mg) = 0.001 gram (g)
micro	μ	10-6	1 microliter (μL) = 0.000001 liter (L)

Scientific Notation

Scientific notation is used to express very large or very small numbers conveniently.

• General Format: $a \times 10^n$ where 1 ≤ a < 10 and n is an integer.

• Example: 0.000073 = $7.3 \times 10^{-5}$

• Example: 945.7 = $9.457 \times 10^{2}$

Significant Figures (SF)

Significant figures are the digits in a measurement that are known with certainty plus one estimated digit. They reflect the precision of a measurement.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros are significant only if there is a decimal point.

Rules for Calculations:

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

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

Exact vs. Measured Numbers

• Exact Numbers: Values known with complete certainty (e.g., counted items, defined quantities such as 1 dozen = 12).

• Measured Numbers: Values obtained by measurement, subject to uncertainty.

Unit Conversions and the Factor-Label Method

Unit conversions use conversion factors, which are ratios derived from equalities between different units.

• Conversion Factor: A fraction that expresses the relationship between two different units (e.g., $\frac{2.54\,\text{cm}}{1\,\text{in}}$).

• To convert units, multiply by the appropriate conversion factor so that units cancel and the desired unit remains.

Example: Convert 184 cm to inches:

• $184\,\text{cm} \times \frac{1\,\text{in}}{2.54\,\text{cm}} = 72.4\,\text{in}$

Temperature Scales and Conversions

Temperature can be measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K). Conversions between these scales use specific formulas:

• $T_{\text{F}} = 1.8 \times T_{\text{C}} + 32$

• $T_{\text{K}} = T_{\text{C}} + 273.15$

Energy and Heat

• Energy: The capacity to do work.

• Heat: A form of energy that flows from hot to cold objects.

• Units: Joule (J), calorie (cal); $1\,\text{cal} = 4.184\,\text{J}$

• Food Calories: 1 Calorie (Cal) = 1 kilocalorie (kcal) = 1000 cal

Specific Heat

Specific heat is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Formula: $\text{Heat} = \text{SH} \times \text{mass} \times \Delta T$

• Where SH = specific heat, mass in grams, $\Delta T$ = change in temperature (°C)

Density and Specific Gravity

Density is the ratio of mass to volume. Specific gravity compares the density of a substance to the density of water.

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

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

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

Summary Table: Key Physical Quantities and Units

Quantity	SI Unit	Metric Unit	Common Equivalents
Mass	kilogram (kg)	gram (g)	1 kg = 2.205 lb; 1000 g = 1 kg; 1 lb = 454 g
Length	meter (m)	meter (m)	1 m = 1.094 yd; 1 in = 2.54 cm; 1 m = 100 cm
Volume	cubic meter (m3)	liter (L)	1 L = 1.057 qt; 1 L = 1000 mL; 1 mL = 1 cm3
Temperature	Kelvin (K)	Celsius (°C)	See conversion formulas above
Energy	joule (J)	calorie (cal)	1 cal = 4.184 J

Example Application: To calculate the density of a cube of lithium metal with mass 0.794 g and volume (0.82 cm × 1.45 cm × 1.25 cm):

• Volume = 0.82 × 1.45 × 1.25 = 1.48625 cm3

• Density = 0.794 g / 1.48625 cm3 = 0.534 g/cm3

Additional info: These notes cover foundational concepts in GOB Chemistry, including matter, measurement, classification, and basic calculations, which are essential for further study in chemistry and related health sciences.