Accuracy and Precision

Understanding Measurement Quality

In chemistry, measurements must be both accurate and precise to ensure reliable results. These concepts are fundamental to laboratory work and data analysis.

• Accuracy: Refers to how close a measured value is to the true or accepted value.

• Precision: Indicates how close repeated measurements are to each other, regardless of their closeness to the true value.

• Example: If a balance gives the same mass reading for a standard weight every time, it is precise. If that reading matches the actual mass, it is also accurate.

Significant Figures

Identifying and Using Significant Figures

Significant figures (sig figs) reflect the precision of a measured or calculated quantity. Proper use of significant figures is essential in reporting scientific data.

• Identifying Significant Figures: 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.

• Rounding: When rounding, retain only as many digits as are significant based on the context or calculation rules.

• Example: 0.00450 has three significant figures (4, 5, and the trailing zero).

Rules for 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.

• Example: (rounded to two significant figures).

Scientific Notation

Expressing Numbers in Scientific Notation

Scientific notation is used to express very large or very small numbers in a compact form.

• Format: , where 1 ≤ a < 10 and n is an integer.

• Example: 0.00056 =

The Metric System and Unit Conversions

Metric Units and Conversions

The metric system is the standard system of measurement in science. Understanding base units and prefixes is essential for converting between units.

• Common Units: meter (m), gram (g), liter (L), second (s), Kelvin (K).

• Prefixes: kilo- (103), centi- (10-2), milli- (10-3), micro- (10-6), nano- (10-9).

• Conversion Example: To convert 2500 mg to grams:

Matter and Its Classification

Physical and Chemical Changes

Matter can undergo physical or chemical changes, each with distinct characteristics.

• Physical Change: Alters the form or appearance but not the composition (e.g., melting ice).

• Chemical Change: Produces new substances with different properties (e.g., rusting iron).

Mixtures and Pure Substances

• Pure Substance: Has a fixed composition (elements or compounds).

• Mixture: Combination of two or more substances not chemically combined.

Homogeneous and Heterogeneous Mixtures

• Homogeneous Mixture: Uniform composition throughout (e.g., saltwater).

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

States of Matter

Properties of Solids, Liquids, and Gases

Matter exists in three primary states, each with unique properties.

Behavior of Gases

Gas Laws and Calculations

Gases have unique behaviors described by several laws. Calculations often involve relationships among pressure, volume, temperature, and amount.

• Boyle's Law: (at constant temperature)

• Charles's Law: (at constant pressure)

• Ideal Gas Law:

• Example: Calculate the volume of 1.0 mol of gas at STP (Standard Temperature and Pressure):

Intermolecular Forces

Types and Effects

Intermolecular forces are attractions between molecules that influence physical properties such as boiling and melting points.

• Types: London dispersion forces, dipole-dipole interactions, hydrogen bonding.

• Example: Water has a high boiling point due to strong hydrogen bonds between molecules.

Temperature Conversions

Celsius and Kelvin Scales

Temperature is measured in Celsius (°C) and Kelvin (K) in chemistry. The Kelvin scale is the SI unit for temperature.

• Conversion Formula:

• Example: 25°C = K

Exothermic and Endothermic Processes

Energy Changes in Chemical Reactions

Chemical reactions involve energy changes, classified as exothermic or endothermic.

• Exothermic: Releases energy (usually as heat) to the surroundings (e.g., combustion).

• Endothermic: Absorbs energy from the surroundings (e.g., photosynthesis).

Density

Calculating Density

Density is a physical property that relates the mass of a substance to its volume.

• Formula:

• Example: If a block has a mass of 50 g and a volume of 20 mL, its density is