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Chemistry Fundamentals

Atomic Structure and Chemistry​

1. Atoms and Atomic Mass​

Atoms (परमाणु):

  • Basic unit of matter
  • Composed of protons, neutrons, electrons
  • Nucleus (protons + neutrons) at center
  • Electrons around nucleus

Atomic Number (Z):

  • Number of protons
  • Defines element
  • Example: Carbon = 6 protons

Mass Number (A):

  • Total protons + neutrons
  • Example: Carbon-12 has 12 total nucleons

Relative Atomic Mass:

  • Based on carbon-12 = 12 standard
  • Indicates mass in atomic mass units (amu)
  • Located on periodic table

Isotopes:

  • Same element, different neutrons
  • Different mass numbers
  • Example: Carbon-12 and Carbon-14

2. Electron Configuration​

Electron Shells (इलेक्ट्रॉन खोल):

  • Organized in energy levels
  • Ring sizes increase outward
  • Shell 1: 2 electrons maximum
  • Shell 2: 8 electrons maximum
  • Shell 3: 18 electrons maximum
  • Valence electrons on outermost shell determine chemical properties

Examples:

  • Hydrogen: 1
  • Carbon: 2,4
  • Oxygen: 2,6
  • Sodium: 2,8,1
  • Chlorine: 2,8,7

Valence Electrons:

  • Outermost shell electrons
  • Determine bonding behavior
  • Example: Carbon (4) can form 4 bonds

Periodic Table​

1. Organization​

Horizontal Rows (Periods):

  • Same number of electron shells
  • Period 1: 1 shell
  • Period 2: 2 shells
  • Period 3: 3 shells

Vertical Columns (Groups):

  • Same number of valence electrons
  • Similar chemical properties
  • Group 1: 1 valence electron (alkali metals)
  • Group 2: 2 valence electrons (alkaline earth metals)
  • Group 17: 7 valence electrons (halogens)
  • Group 18: 8 valence electrons (noble gases)

2. Element Classifications​

Metals (धातु):

  • Conduct electricity and heat
  • Malleable (can be shaped)
  • Ductile (drawn to wire)
  • Lose electrons (form cations)
  • Example: Iron, copper, aluminum

Non-metals (अधातु):

  • Poor conductors
  • Brittle (solid)
  • Gain electrons (form anions)
  • Example: Oxygen, nitrogen, sulfur

Metalloids:

  • Properties between metals and non-metals
  • Semiconductors
  • Example: Silicon, germanium

Chemical Bonding​

1. Ionic Bonding (आयनिक बंधन)​

Definition:

  • Electrons transferred from metal to non-metal
  • Forms cations (positive) and anions (negative)
  • Electrostatic attraction

Process:

  • Metal loses electrons (oxidation)
  • Non-metal gains electrons (reduction)
  • Transfer of electrons
  • Opposite charges attract

Examples:

  • NaCl: Na loses 1e⁻ to Cl

    • Forms Na⁺ and Cl⁻
    • Electrostatic attraction
    • Ionic compound

  • MgO: Mg loses 2e⁻ to O

    • Forms Mg²⁺ and O²⁻
    • Ionic compound

Properties:

  • Crystalline solids (usually)
  • High melting/boiling points
  • Conducted electricity when molten/dissolved
  • Typically soluble in water

2. Covalent Bonding (सहसंयोजक बंधन)​

Definition:

  • Electrons shared between atoms
  • Both attracted to shared electrons
  • Occurs between non-metals

Single Bond:

  • 2 electrons shared
  • Example: H-H in hydrogen
  • Example: C-C in ethane

Double Bond:

  • 4 electrons shared
  • Stronger than single bond
  • Example: C=C in ethene
  • Example: O=O in oxygen

Triple Bond:

  • 6 electrons shared
  • Example: N≡N in nitrogen
  • Strongest bond

Covalent Compounds:

  • Molecules (discrete units)
  • Low melting/boiling points
  • Not conductive
  • Example: Methane (CH₄), water (H₂O)

Polar Covalent:

  • Unequal sharing due to electronegativity difference
  • Partial charges (δ+ and δ-)
  • Example: H-Cl, H-O bonds

Non-polar Covalent:

  • Equal sharing (similar electronegativity)
  • Example: C-C, H-H bonds

3. Metallic Bonding (धात्विक बंधन)​

Definition:

  • Metal atoms lose valence electrons
  • Electrons delocalized ("sea of electrons")
  • Cations held in electron sea
  • Non-directional

Properties:

  • Conduct electricity (mobile electrons)
  • Conduct heat
  • Malleable (atoms can shift)
  • Ductile (can be drawn to wire)
  • Lustrous (reflect light)
  • Example: Cu, Al, Fe

Chemical Equations and Reactions​

1. Writing Chemical Equations (रासायनिक समीकरण)​

Reactants → Products

Example:

  • H₂ + Cl₂ → 2HCl
  • Unbalanced: H₂ + Cl₂ → HCl
  • Balanced: H₂ + Cl₂ → 2HCl
  • Coefficients show mole ratios

States of Matter:

  • (s) = solid
  • (l) = liquid
  • (g) = gas
  • (aq) = aqueous (dissolved)

Example with States:

  • 2H₂(g) + O₂(g) → 2H₂O(l)

2. Types of Chemical Reactions​

Synthesis (संश्लेषण):

  • A + B → AB
  • Elements/compounds combine
  • Example: 2H₂ + O₂ → 2H₂O

Decomposition (विघटन):

  • AB → A + B
  • Compound breaks apart
  • Requires energy
  • Example: 2H₂O₂ → 2H₂O + O₂

Single Displacement:

  • A + BC → AC + B
  • More active replaces less active
  • Example: Zn + 2HCl → ZnCl₂ + H₂

Double Displacement:

  • AB + CD → AD + CB
  • Ions exchange partners
  • Often produces precipitate
  • Example: AgNO₃ + NaCl → AgCl↓ + NaNO₃

Combustion (दहन):

  • Fuel + O₂ → CO₂ + H₂O + Energy
  • Burning in oxygen
  • Exothermic (releases energy)
  • Example: CH₄ + 2O₂ → CO₂ + 2H₂O

Acid-Base Neutralization:

  • Acid + Base → Salt + Water
  • H⁺ + OH⁻ → H₂O
  • Example: HCl + NaOH → NaCl + H₂O

3. Balancing Equations​

Steps:

  1. Write unbalanced equation
  2. Count atoms of each element on both sides
  3. Add coefficients to balance (multiply compounds)
  4. Check all atoms balanced
  5. Use smallest whole numbers

Example:

  • Unbalanced: Fe + O₂ → Fe₂O₃
  • Count: Fe: 1 vs 2; O: 2 vs 3 (not balanced)
  • Add coefficients: 4Fe + 3O₂ → 2Fe₂O₃
  • Check: Fe: 4 vs 4 ✓; O: 6 vs 6 ✓ (balanced)

Stoichiometry (रासायनिक मात्रा)​

1. Molar Mass and Moles​

Molar Mass:

  • Mass of one mole of substance
  • Sum of atomic masses
  • Units: g/mol
  • Example: H₂O = 2(1) + 16 = 18 g/mol

Mole (मोल):

  • 6.02 × 10²³ particles (Avogadro's number)
  • Converts between atoms/molecules and grams
  • Allows use of balanced equations

Conversions:

  • Grams to moles: n = mass ÷ molar mass
  • Moles to grams: mass = n × molar mass
  • Moles to particles: # = n × 6.02 × 10²³

2. Stoichiometric Calculations​

Using Balanced Equations:

  • Mole ratios from coefficients
  • Example: 2H₂ + O₂ → 2H₂O
  • Mole ratio H₂:O₂ = 2:1
  • If 4 mol H₂, then 2 mol O₂ needed

Mass to Mass Calculation:

  1. Convert grams to moles (mass ÷ molar mass)
  2. Use mole ratio from equation
  3. Convert moles to grams (mol × molar mass)

Example:

  • 2H₂ + O₂ → 2H₂O
  • How many grams of water from 4g H₂?
  • 4g H₂ ÷ 2 g/mol = 2 mol H₂
  • Mole ratio 2H₂:2H₂O = 1:1
  • 2 mol H₂ produces 2 mol H₂O
  • 2 mol × 18 g/mol = 36g H₂O

3. Limiting Reactant and Yield​

Limiting Reactant (सीमित अभिकारक):

  • Reactant that runs out first
  • Determines maximum product
  • Other reactants in excess

Theoretical Yield:

  • Maximum product possible
  • Assuming 100% conversion
  • From limiting reactant

Percent Yield:

  • Actual ÷ Theoretical × 100
  • Shows efficiency
  • Always ≤100%
  • Accounts for losses

Acids, Bases, and pH​

1. Acids and Bases​

Acids (अम्ल):

  • Produce H⁺ ions (or donate protons)
  • Taste sour
  • pH < 7
  • Example: HCl, H₂SO₄, acetic acid

Bases (क्षार):

  • Produce OH⁻ ions (or accept protons)
  • Taste bitter
  • pH > 7
  • Example: NaOH, ammonia

Neutralization:

  • Acid + Base → Salt + Water
  • H⁺ + OH⁻ → H₂O
  • Exothermic

2. pH Scale​

pH = -log[H⁺]

  • Range: 0-14
  • pH 7: Neutral
  • pH < 7: Acidic
  • pH > 7: Basic/Alkaline
  • pH 0: Most acidic
  • pH 14: Most basic

Summary​

Chemistry fundamentals:

  • Atoms: Basic building blocks of matter
  • Bonding: Atoms combine through ionic, covalent, or metallic bonds
  • Reactions: Chemical changes from combination or breaking of bonds
  • Equations: Represent reactions with balanced formulas
  • Stoichiometry: Quantitative relationships in reactions
  • Acids/Bases: Important categories of compounds

Chemistry explains composition, structure, and transformations of matter.