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 knowledge.

• Property: A characteristic useful for identifying a substance (e.g., size, color, temperature, chemical composition, reactivity).

• Physical Change: Alters the form or appearance of matter without changing its chemical identity (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 pure substances or mixtures.

• Pure Substance: Uniform chemical composition throughout (e.g., elements, compounds).

• Mixture: Blend of two or more substances, each retaining its 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 (e.g., oxygen, gold).

• Compound: Pure substance that can be broken down into simpler substances by chemical reactions (e.g., water, carbon dioxide).

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-letter abbreviations (first letter capitalized, second lowercase).

• Some symbols are based on Latin names (e.g., Na for sodium from natrium).

• Chemical Formula: Notation using element symbols and subscripts to indicate the number of atoms (e.g., H2O).

• Elements essential for life include C, H, O, N, and others (macronutrients and micronutrients).

Chemical Reactions: Examples of Chemical Change

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

• Chemical Reaction: Process in which the identity and composition of substances change.

• 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.

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 quantity for measurement (e.g., kilogram, meter, liter, kelvin, second).

• SI Units: International System of Units (e.g., kg for mass, m for length, K for temperature).

• Metric Units: Gram (g), liter (L), Celsius (°C).

• Derived Units: Formed from base units (e.g., m/s for speed, g/cm3 for density).

• Prefixes: Modify unit size (e.g., milli-, micro-, kilo-).

• Scientific Notation: Expresses numbers as , where and is an integer.

Example: 0.000000120 m = m = 120 nm.

Measuring Mass, Length, and Volume

Understanding the difference between mass and weight and knowing the appropriate units is essential in chemistry.

• Mass: Amount of matter in an object (measured in grams or kilograms).

• Weight: Gravitational force on an object (varies with location).

• Length: Measured in meters (m), centimeters (cm), millimeters (mm).

• Volume: Amount of space occupied (measured in liters (L), milliliters (mL), or cubic centimeters (cm3)).

Example: 1 L = 1000 mL; 1 mL = 1 cm3.

Measurement and Significant Figures

All measurements have some uncertainty, and significant figures reflect the precision of a measurement.

• 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 have unlimited 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, based on the precision of the data.

• 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 laboratory and medical calculations. The factor-label (dimensional analysis) method is commonly used.

• Conversion Factor: Ratio expressing the relationship between two units (e.g., 1 in = 2.54 cm).

• Factor-Label Method: Set up calculations so that units cancel, leaving only 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.

  4. Solve, checking that units cancel appropriately.

  5. Make a ballpark estimate to check reasonableness.

Example: To convert 21.5 in to cm:

Temperature, Heat, and Energy

Temperature and heat are related but distinct concepts. Temperature is measured in different scales, and energy changes accompany all chemical reactions.

• 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:

• Specific Heat: Amount of heat needed to raise the temperature of 1 g of a substance by 1°C.

• Heat Calculation: Where = heat (J or cal), = mass (g), = specific heat (J/g·°C or cal/g·°C), = temperature change (°C).

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

Density and Specific Gravity

Density relates mass to volume and is a key property for identifying substances. Specific gravity compares the density of a substance to that of water.

• Density: (units: g/cm3 for solids, g/mL for liquids).

• Specific Gravity: (unitless).

• Substances with density less than water float; those with greater density sink.

• Hydrometers and urinometers are used to measure specific gravity in liquids.

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

Tables

Key tables from the chapter (recreated and summarized):

Additional info: Some table entries and values are inferred or summarized for clarity and completeness.