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.

• Matter: The physical material of the universe; anything with mass and volume.

• Scientific Method: A systematic approach involving observation, hypothesis, and experimentation to expand scientific knowledge.

• Property: A characteristic used to identify substances, such as size, color, temperature, chemical composition, and reactivity.

• Physical Change: Alters the form or appearance of matter without changing its chemical identity (e.g., melting, dissolving).

• Chemical Change: Results in the formation of new substances with different properties (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: The transformation 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: Combination of two or more substances, each retaining its identity.

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

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

• Element: A fundamental substance that cannot be broken down chemically (e.g., oxygen).

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

Example: Vanilla ice cream is a homogeneous mixture; sugar is a compound.

Chemical Elements and Symbols

There are 118 known elements, 91 of which occur naturally. Chemists use symbols to represent elements:

• Symbols consist of one or two letters (first capitalized, second lowercase).

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

• Elements combine to form compounds, represented by chemical formulas (e.g., H2O).

• Subscripts in formulas indicate the number of atoms of each element.

Chemical Reactions: Examples of Chemical Change

Chemical reactions are processes in which substances (reactants) are transformed into new substances (products):

• Chemical Reaction: Change in identity and composition of substances.

• Reactant: Starting material in a reaction.

• Product: Substance formed as a result of the reaction.

• Reactions are represented with reactants on the left, products on the right, and an arrow indicating the direction of change.

Example: Electrolysis of water produces hydrogen and oxygen gases.

Physical Quantities: Units and Scientific Notation

Physical properties are measured using standardized units. The SI (International System of Units) is the standard in science:

• Mass: kilogram (kg)

• Length: meter (m)

• Volume: cubic meter (m3), but liter (L) is commonly used

• Temperature: kelvin (K)

• Time: second (s)

Prefixes are used to indicate multiples or fractions of units (e.g., milli-, micro-, kilo-).

Scientific Notation: Used to express very large or small numbers as a product of a number between 1 and 10 and a power of 10.

• Example:

Measuring Mass, Length, and Volume

Understanding the difference between mass and weight is crucial:

• 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, centimeters, or millimeters.

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

Measurement and Significant Figures

All measurements have some uncertainty, reflected in the number of significant figures reported:

• 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 may or may not be significant.

• Exact Numbers: Have an infinite number of significant figures (e.g., 1 foot = 12 inches).

Example: 0.0076 mL has two significant figures; 3400 kg may have two, three, or four significant figures depending on context.

Rounding Off Numbers

When performing calculations, the number of significant figures in the result is determined by the precision of the input values:

• Multiplication/Division: Result has as many significant figures as the least precise value.

• Addition/Subtraction: Result has as many decimal places as the least precise value.

• 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: 125.884 rounded to the nearest whole number is 126.

Problem Solving: Unit Conversions and Estimating Answers

Unit conversions are essential in chemistry. The factor-label (dimensional analysis) method is used to convert between units:

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

• Set up calculations so that unwanted units cancel, leaving the desired unit.

• Always check your answer with a ballpark estimate to ensure it is reasonable.

Example: To convert 21.5 inches to centimeters:

Temperature, Heat, and Energy

All chemical reactions involve energy changes, often in the form of heat. Temperature is a measure of the average kinetic energy of particles in a substance.

• Energy: The capacity to do work or supply heat (measured in joules, J).

• Temperature Scales:

  • Celsius (°C): Freezing point of water = 0°C, boiling point = 100°C.

  • Kelvin (K): Absolute zero = 0 K;

  • Fahrenheit (°F):

• Calorie (cal): Amount of heat needed to raise 1 g of water by 1°C.

• Specific Heat (c): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

Formula for heat:

where = heat (cal or J), = mass (g), = specific heat, = temperature change (°C).

Example: Heating 95 kg of water by 25°C requires

Density and Specific Gravity

Density is a key physical property relating mass and volume:

• Density (d):

• Units: g/cm3 for solids, g/mL for liquids.

• Substances with lower density float on those with higher density.

• Specific Gravity: Ratio of the density of a substance to the density of water at the same temperature (unitless).

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

Summary Table: Key Physical Quantities and Units

Additional info: Some tables and figures referenced in the original text (e.g., Table 1.1, 1.3, 1.4, 1.5, 1.6, 1.7, 1.8, 1.9, 1.10, 1.11) are not fully reproduced here, but their main content has been summarized or inferred based on standard chemistry knowledge.