Basic Concepts of Chemistry

1. Matter

Types of Matter:
- Pure Substances: Contain the same type of particles.
  - Examples:
    - Iron sheet (contains one type of particle)
    - Sodium chloride (NaCl)
- Impure Substances (Mixtures): Contain different types of particles.
  - Examples:
    - Salt solution (contains salt and water)
    - Salad (contains various ingredients)
Key Exam Question: Why are elements and compounds called pure substances?
- Answer: They are made up of one type of particle.
Mixtures: Impure substances made up of different types of particles.
- Important examples include: air, sea water, solutions, suspensions, colloids, and alloys like steel.

Physical Properties: Can be observed without changing the identity of the substance.
- Mnemonic: "Moving Vent Should Come Late"
  - M = Mass
  - V = Volume
  - S = Shape
  - C = Color
  - L = Length
Chemical Properties: Observed when a substance undergoes a change in state.
- Mnemonic: "Moving Vent Should Come Late" (for chemical reactions)
  - A = Acid
  - S = Air
  - B = Base
  - W = Water
  - C = Chemical reactions

Definition: Anything that can be measured.
Types:
- Base Physical Quantities:
  - Length, Time, Mass, Temperature, Amount of Substance, Light Intensity, Electric Current.
- Derived Physical Quantities: Derived from base quantities (e.g., Speed = Length/Time).
International System of Units (SI Units): Established in 1960 for base quantities.
- Mnemonic: "Lisa Mem Turns to a Left Age"
  - L = Length (m)
  - M = Mass (kg)
  - T = Time (s)
  - T = Temperature (K)
  - S = Amount of Substance (mol)
  - L = Light Intensity (cd)
  - E = Electric Current (A)
Difference between Mass and Weight:
- Mass: Amount of matter in an object (SI unit: kg)
- Weight: Force of gravity acting on an object (SI unit: N).

Definition: Mnemonics added to units to show multiples or fractions.
- Example: 2 kg (kilo is the prefix).
Important Prefixes:
- Positive:
  - D = Deca
  - H = Hecto
  - K = Kilo
  - M = Mega
  - G = Giga
  - T = Tera
- Negative:
  - D = Deci
  - C = Centi
  - M = Milli
  - U = Micro
  - N = Nano
  - P = Pico

Definition: A method for writing very large or very small numbers.
- Example: Mass of the earth = 6 x 10^24 kg
- Diameter of nucleus = 1.7 x 10^-15 m

Definition: Certain and important digits in a measurement.
Identifying Significant Figures:
- Non-decimal numbers: Count from the first non-zero digit to the last non-zero digit.
- Decimal numbers: Count from the first non-zero digit to the last digit.

Accuracy: How close the measurement is to the actual value.
Precision: How close multiple measurements are to each other.
Examples:
- Accurate but not precise = One measurement matches actual value, others don’t.
- Precise but not accurate = All measurements are close to each other but not to the actual value.

Definition: The process of converting one set of units to another.
- Example: 3045 m to km (3045 m / 1000 = 3.045 km).
- Example: 35 °C to K (35 + 273 = 308 K).

Laws:
1. Law of Conservation of Mass: Mass cannot be created or destroyed in a chemical reaction.
2. Law of Definite Proportions: Elements combine in fixed ratios by mass.
3. Law of Multiple Proportions: When two elements form different compounds, the mass ratio of the elements is a whole number.
4. Gay-Lussac's Law: Pressure is directly proportional to temperature at constant volume.
5. Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

Definition of Mole: A counting unit similar to a dozen.
Avogadro's Number: 6.022 x 10^23 particles (atoms, molecules).
Molar Mass: Mass of one mole of a substance (e.g., 12 g of carbon = 1 mole).

Molarity (M): Moles of solute per liter of solution.
- Formula: M = moles of solute / volume of solution (L).
Molality (m): Moles of solute per kg of solvent.
- Formula: m = moles of solute / mass of solvent (kg).

These notes summarize the essential concepts from the first chapter of chemistry. Remember to review them for a solid understanding.