Introduction to Chemistry

Definition and Scope

Chemistry is the study of matter, its properties, and the changes it undergoes. Matter is anything that has mass and occupies space.

• Atoms: The basic unit of matter.

• Molecules: Two or more atoms bonded together.

• Compounds: Molecules composed of two or more different types of atoms.

• Mixtures: Physical combinations of elements and/or compounds.

Elements and Compounds

Elements

• An element is a unique kind of atom; it may not be broken down into simpler substances by chemical means.

• Elements are represented by symbols (one or two letters, first always capitalized).

• Oxygen is an important element in many compounds.

Law of Constant Composition

Compounds have a definite composition, meaning the relative number of atoms of each element in a compound is the same in every sample.

Properties of Matter

Physical and Chemical Properties

• Physical properties: Can be observed without changing the substance into another substance (e.g., density, boiling point, color).

• Chemical properties: Can only be seen when a substance is changed into another substance (e.g., flammability, reactivity).

Intensive and Extensive Properties

• Intensive properties: Independent of the amount of substance (e.g., density, boiling point, color).

• Extensive properties: Depend on the amount of substance (e.g., mass, volume, energy).

Physical and Chemical Changes

• Physical changes: Changes in matter that do not change the substance itself (e.g., changes of state, temperature, volume).

• Chemical changes: Result in a new substance (e.g., combustion, oxidation, decomposition).

Separation of Mixtures

Methods of Separation

• Filtration: Separates solids from liquids (e.g., sand from water).

• Distillation: Uses differences in boiling points to separate a liquid homogeneous mixture into its components.

• Chromatography: Separates substances based on their ability to adhere to a solid surface.

Energy and Work

Definitions

• Energy: The capacity to do work or transfer heat.

• Work: The energy transferred when a force is applied to an object causing movement.

• Force: Any push or pull on an object.

Kinetic and Potential Energy

• Kinetic energy: Energy of motion, depends on mass and velocity.

• Formula:

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

• Formula:

Units of Measurement

Quantitative Measurements

• Measured using SI units (International System of Units).

• Examples: meter (length), kilogram (mass), second (time), kelvin (temperature), mole (amount of substance).

• Luminous intensity: candela (cd).

SI Prefixes

SI prefixes are used to indicate multiples or fractions of units.

Temperature Scales

• Kelvin (K): The SI unit for temperature. Absolute zero (0 K) is the lowest possible temperature.

• Conversion formulas:

  • Kelvin to Celsius:

  • Fahrenheit to Celsius:

  • Celsius to Fahrenheit:

Volume and Mass Units

• 1 liter (L) = 1 cubic decimeter (dm3)

• 1 milliliter (mL) = 1 cubic centimeter (cm3)

• 1 kilogram (kg) = 2.20 pounds (lbs)

• 1 meter (m) = 1.09 yards

Laboratory Measurement Tools

Common Volumetric Equipment

• Graduated cylinder: Used to measure liquid volumes accurately.

• Syringe: Used to deliver variable volumes.

• Burette: Used for titration to deliver precise volumes.

• Pipette: Used to deliver a specific volume.

• Volumetric flask: Used to prepare solutions of precise volume.

Density

Definition and Formula

• Density is the mass per unit volume of a substance.

• Formula:

Common Densities

Measurement, Precision, and Accuracy

Uncertainty and Significant Figures

• All measurements have some degree of uncertainty.

• Exact numbers: Known exactly (e.g., counting numbers, defined quantities).

• Inexact numbers: Measured using scientific instruments; always have some uncertainty.

• Digit of uncertainty: The last digit measured is reliable but not exact.

Precision and Accuracy

• Precision: How closely repeated measurements agree with each other.

• Accuracy: How close a measurement is to the true value.

Significant Figures (Sig Figs)

• Rules for determining significant figures:

  1. All nonzero digits are significant.

  2. Zeros between nonzero digits are significant.

  3. Zeros at the beginning are not significant.

  4. Zeros at the end are significant if there is a decimal point.

• For addition/subtraction: The result has the same number of decimal places as the measurement with the fewest decimal places.

• For multiplication/division: The result has the same number of significant figures as the measurement with the fewest significant figures.

Energy Units and Conversions

• Joule (J): SI unit of energy. 1 J = energy of 2 kg moving at 1 m/s.

• Calorie (cal): 1 cal = 4.18 J.

• 1 nutritional Calorie (Cal) = 1 kcal = 1000 cal.

Examples

• Example of density calculation: If a sample has a mass of 10 g and a volume of 2 cm3, its density is .

• Example of significant figures: 0.00450 has three significant figures (the leading zeros are not significant).

Additional info: Some context and explanations have been expanded for clarity and completeness based on standard General Chemistry curriculum.