Chapter 2: Chemistry and Measurements

Metric and SI Units of Measurement

The metric system and the International System of Units (SI) are standard systems used in chemistry for measuring quantities such as length, mass, volume, and temperature.

• Metric Units: Commonly used units include meter (m) for length, gram (g) for mass, and liter (L) for volume.

• SI Units: The SI system uses the meter (m), kilogram (kg), and cubic meter (m3), among others.

• Prefixes: Prefixes such as kilo-, centi-, and milli- indicate multiples or fractions of units (e.g., 1 kilometer = 1000 meters).

Significant Figures (Sig Figs)

Significant figures reflect the precision of a measured or calculated quantity.

• Identifying Sig Figs: 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.

• Measuring with Sig Figs: When measuring with a ruler or graduated cylinder, record all certain digits plus one estimated digit.

• Scientific Notation: Express numbers as a product of a coefficient (with correct sig figs) and a power of ten (e.g., 3.20 × 102).

Calculations with Significant Figures

• 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.

Conversions and Prefixes

Unit conversions use conversion factors based on metric prefixes or relationships between units.

• Example: To convert 250 mg to grams: 250 mg × (1 g / 1000 mg) = 0.250 g

Density and Volume Displacement

Density is a physical property that relates mass and volume.

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

• Units: Commonly g/mL or g/cm3

• Volume Displacement: Used to measure the volume of irregular objects by the amount of water they displace.

• Example: If a 10.0 g object displaces 2.5 mL of water, its density is $Density = \frac{10.0\ g}{2.5\ mL} = 4.0\ g/mL$

Specific Gravity

Specific gravity compares the density of a substance to the density of water (1.00 g/mL at 4°C).

• Formula: $Specific\ gravity = \frac{density\ of\ sample}{density\ of\ water}$

• No units: Specific gravity is a ratio and therefore unitless.

Key Equations to Memorize

• $Density = \frac{mass}{volume}$

• $D = \frac{m}{v}$

• $Specific\ gravity = \frac{density\ of\ sample}{density\ of\ water}$

Chapter 3: Matter and Energy

Matter and Its Classification

Matter is anything that has mass and occupies space. It can be classified based on composition and properties.

• Pure Substances: Have a fixed composition (elements or compounds).

• Mixtures: Physical combinations of two or more substances.

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

• Heterogeneous Mixtures: Non-uniform composition (e.g., salad, sand and iron filings).

Physical vs. Chemical Properties and Changes

• Physical Properties: Observed without changing the substance (e.g., color, melting point, density).

• Chemical Properties: Describe how a substance reacts (e.g., flammability, reactivity).

• Physical Change: Change in state or appearance without changing composition (e.g., melting, boiling).

• Chemical Change: Produces new substances (e.g., rusting, burning).

States of Matter

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume, no definite shape; particles can move past each other.

• Gas: No definite shape or volume; particles are far apart and move freely.

Temperature Conversions

• Celsius to Kelvin: $K = ^\circ C + 273.15$

• Celsius to Fahrenheit: $^\circ F = (^\circ C \times \frac{9}{5}) + 32$

• Fahrenheit to Celsius: $^\circ C = (^\circ F - 32) \times \frac{5}{9}$

Energy: Kinetic and Potential

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Unit: Joule (J) or calorie (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: $Specific\ Heat = \frac{heat}{grams \times (Temperature_2 - Temperature_1)}$

• Alternate: $SH = \frac{heat}{grams \times \Delta T}$

• Units: $\frac{cal\ (or\ J)}{g\ ^\circ C}$

• Example: If 100 cal of heat raises 20 g of water by 5°C, $SH = \frac{100\ cal}{20\ g \times 5^\circ C} = 1.0\ cal/g^\circ C$

Phase Changes and Heating/Cooling Curves

• Melting Point/Freezing Point: Temperature at which a substance changes between solid and liquid.

• Boiling Point/Condensation Point: Temperature at which a substance changes between liquid and gas.

• Sublimation: Solid to gas (e.g., dry ice).

• Deposition: Gas to solid.

• Heating/Cooling Curves: Graphs showing temperature changes as heat is added or removed, indicating phase changes.

Table: Classification of Matter

Additional info: This guide expands on the listed topics with definitions, examples, and equations to provide a comprehensive review for exam preparation.