Chapter 1
Chapter 1: Matter and Measurement
Page 2: What is Chemistry? Definition of Chemistry: The study of matter, its composition, properties, and transformations. Matter: Anything that has mass and occupies volume. States of Matter:
Solid State:
Definite volume
Maintains shape regardless of the container
Particles are close together in a regular pattern.
Liquid State:
Definite volume
Assumes the shape of its container
Particles are close but move past one another.
Gas State:
No definite volume, takes the shape of the container
Particles are far apart and move randomly.
Page 3: Properties of Matter
Physical Properties: Observable or measurable without changing the composition. Examples include:
Boiling point
Melting point
Solubility
Color
Odor
State of matter
Physical Change: Alters a material without changing its composition.
Chemical Properties: Indicate how a substance can convert into another.
Chemical Change: A chemical reaction converting one substance into another.
Page 4: Classification of Matter
Pure Substance:
Composed of a single component (atom or molecule)
Constant composition, not dependent on sample size
Cannot be broken down into other pure substances by physical change.
Mixtures:
Composed of multiple components
Varying composition across solid, liquid, and gas
Can be separated by physical changes (e.g., sugar dissolved in water).
Page 5: Elements and Compounds
Elements: Pure substances that cannot be chemically broken down.
Compounds: Pure substances formed by chemically joining two or more elements.
Page 6: Measurement Basics
Every measurement consists of a number and a unit.
Base units in the metric system:
Length: meter (m)
Mass: gram (g)
Prefixes indicate size relative to the base unit (e.g., kilo-, centi-).
Page 7: Metric Prefixes and Examples Conversions:
1 m = 10 dm = 100 cm = 1,000 mm
1 kg = 1,000 g
1 g = 1,000 mg
1 L = 1,000 mL
Exact vs. Inexact Numbers:
Exact numbers come from counting (10 fingers) or definitions (1 m = 100 cm).
Inexact numbers derive from measurements and have uncertainty (e.g., 15.3 cm).
Page 8: Significant Figures
Definition: All digits in a measurement including one estimated digit.
Significant Figures:
All nonzero digits are significant.
Example: 65.2 g (3 significant figures).
Rules for Zeros:
Count as significant between nonzero digits (29.05 g).
Count as significant at the end of a decimal number (3.7500 cm).
Page 9: Zero Rules Continued and Operations
Non-Significant Zero Rules:
Leading zeros do not count (0.00245 mg).
Trailing zeros without decimals do not count (2570 m).
Multiplication/Division Rule: Result has the same number of significant figures as the least precise number.
Page 10: Rounding Rules
Rounding Off: If the first digit dropped is:
0-4: Drop it.
5-9: Round up the last digit.
Examples:
61.2537 rounded to 2 decimal places is 61.25.
Page 11: Addition and Subtraction Rules
Rule for Addition/Subtraction: Result has the same number of decimal places as the number with the fewest decimal places.
Example: 10.11 kg - 3.6 kg = 6.5 kg (1 decimal place).
Page 12: Scientific Notation
Format: Written as Coefficient (between 1 and 10) and Exponent (whole number).
Conversion Steps:
Move the decimal to get a number between 1 and 10.
Multiply by 10 raised to x (number of decimal places moved).
Page 13: Conversion Factors
Conversion: Original quantity multiplied by conversion factor gives desired quantity.
Example: 2.205 lb = 1 kg can be used to convert units.
Page 14: Unit Cancellation in Conversions
Ensure units cancel: For example, converting 130 lb to kg requires that lb cancels out.
Page 15: Solving Conversion Problems
Example: Convert 325 mg of aspirin to grams.
Identify original (325 mg) and desired (g) quantities.
Page 16: Using Multiple Conversion Factors
Temperature: Method to convert quantities involves multiple factors to achieve the desired unit.
Page 17: Measuring Temperature
Temperature Scales: Fahrenheit (°F), Celsius (°C), Kelvin (K).
°C and K are equivalent but offset by 273.