Answers

1. & 2. Mixtures: Definition and Types

• Definition: A mixture is a combination of two or more substances where each substance retains its own chemical identity and properties.
• Homogeneous vs. Heterogeneous:
  • Homogeneous: Uniform composition and properties throughout. You cannot see the individual components (e.g., air, saltwater, brass).
  • Heterogeneous: Non-uniform composition. You can see the different parts (e.g., sand and water, pizza, granite).

3. Properties of Elements

• Physical Properties: (Observed without changing composition)
  • Example for Sodium (Na): Silvery-white metal, soft, low melting point (97.8°C), good conductor of heat and electricity.
• Chemical Properties: (Describes reactivity)
  • Example for Sodium (Na): Highly reactive with water (2Na + 2H₂O → 2NaOH + H₂), reacts violently with chlorine to form table salt (2Na + Cl₂ → 2NaCl).

4. Isotopes: Math related to Isotopes → Atomic Mass

This involves calculating the Average Atomic Mass of an element, which is a weighted average based on the mass and natural abundance of each isotope.

Formula: Average Atomic Mass = (Mass of Isotope₁ × Abundance₁) + (Mass of Isotope₂ × Abundance₂) + ...

Example Calculation: Let's use the example of Chlorine from your textbook:

• Isotope Cl-35: Mass = 34.97 amu, Abundance = 75.76% (or 0.7576)
• Isotope Cl-37: Mass = 36.97 amu, Abundance = 24.24% (or 0.2424)

Average Mass = (34.97 amu × 0.7576) + (36.97 amu × 0.2424) Average Mass = 26.50 amu + 8.96 amu Average Mass = **35.46 amu**

This is the value you see for chlorine on the periodic table.

5. Percentage & Mass Calculation: An element will be given with %. Calculate how much (in grams) of a certain element is present in a compound.

This is called finding the Percent Composition by Mass.

Formula: % Mass of Element = (Mass of Element in 1 mol of Compound / Molar Mass of Compound) × 100%

Your Example: "How much H₂ is present (both % and gram) in H₂O?"

• Step 1: Find the molar mass of H₂O.

  • H: 2 atoms × 1.01 g/mol = 2.02 g/mol
  • O: 1 atom × 16.00 g/mol = 16.00 g/mol
  • Molar Mass of H₂O = 2.02 + 16.00 = 18.02 g/mol

• Step 2: Find the mass contributed by H.

  • This is the "Mass of Element" part. In H₂O, the mass from hydrogen is 2.02 g.

• Step 3: Calculate the Percentage.

  • % H = (2.02 g / 18.02 g) × 100%
  • % H = **11.2%**

• Step 4: Calculate the mass in a given sample.

  • The percentage means that in any sample of water, 11.2% of its mass is hydrogen.
  • Therefore, in 100 grams of water, there are 11.2 grams of hydrogen.

6. Chemical Formula Names: Write name from formula.

The rules depend on whether the compound is Ionic (metal + nonmetal) or Molecular (nonmetals only).

1. Ionic Compounds (e.g., NaCl, CaO, MgF₂)

• Rule: Name the metal (cation) first, then the nonmetal (anion) with its ending changed to -ide.
• Examples:
  • NaCl = Sodium chloride
  • CaO = Calcium oxide
  • MgF₂ = Magnesium fluoride
  • Al₂O₃ = Aluminum oxide

2. Molecular (Covalent) Compounds (e.g., N₂O₅, CO₂)

• Rule: Use prefixes to indicate the number of each atom. The second element ends in -ide.
  • Prefixes: 1 = mono-, 2 = di-, 3 = tri-, 4 = tetra-, 5 = penta-, 6 = hexa-
• Examples:
  • N₂O₅ = Dinitrogen pentoxide
  • CO₂ = Carbon dioxide (the "mono-" prefix is dropped for the first element)
  • P₄O₁₀ = Tetraphosphorus decoxide

7., 9., 10. Equilibrium

• Equilibrium Reaction: A reversible reaction where the rates of the forward and reverse reactions are equal, and the concentrations of reactants and products remain constant.

Kc

Equilibrium constant expressed in terms of the molar concentration of reactants and products.

Formula: Kc = ([C]^c [D]^d) / ([A]^a [B]^b)

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Kp

Equilibrium constant expressed in terms of the partial pressures of gaseous reactants and products.

Formula: Kp = (PC^c PD^d) / (PA^a PB^b)

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Relation between Kp and Kc

Kp = Kc (RT)^Δn where Δn = (moles of gaseous products - moles of gaseous reactants), R = 0.0821 L·atm·mol⁻¹·K⁻¹, T = Temperature in Kelvin.

• Kc vs. Kp:

  • Kc: Equilibrium constant expressed using concentrations (mol/L).
  • Kp: Equilibrium constant expressed using partial pressures (atm).

• Equilibrium Shift (Le Châtelier's Principle): If a stress is applied to a system at equilibrium, the system shifts to relieve that stress.

Quick Guide: Predicting Equilibrium Shift

8. Empirical & Molecular Formulas

• Steps to find Empirical Formula from % Composition:
  1. Assume 100g sample; % becomes grams.
  2. Convert grams to moles for each element.
  3. Divide all mole values by the smallest number of moles.
  4. If needed, multiply by a small integer or float to get whole numbers.
• Molecular Formula = (Empirical Formula)ₙ, where n = (Molar Mass) / (Empirical Formula Mass).

11. Acid-Base Theories

1. Arrhenius (Your PDF): Acid produces H⁺ ions in water; base produces OH⁻ ions.
2. Brønsted-Lowry (Your PDF): Acid is a proton (H⁺) donor; base is a proton acceptor.
3. Lewis (Your PDF): Acid is an electron pair acceptor; base is an electron pair donor.

12. pH Calculations

• From [H⁺]: pH = -log[H⁺]
• From [OH⁻]: pOH = -log[OH⁻], then pH = 14 - pOH
• Example: What is the pH of a solution with [OH⁻] = 1.0 × 10⁻³ M?
  • pOH = -log(1.0 × 10⁻³) = 3.00
  • pH = 14.00 - 3.00 = 11.00

13. Chapter 1 Review: Everything from the 1st chapter.

Valence Electrons

• Definition: Electrons in the outermost shell, responsible for bonding.
• How to determine: Use Group Number (except transition metals).
• Example:
  • Na (Z=11) → 1 valence e⁻ (Group 1).
  • O (Z=8) → 6 valence e⁻ (Group 16).
  • Cl (Z=17) → 7 valence e⁻ (Group 17).
• Trend:
  • Across a Period → Increases from 1 → 8.
  • Down a Group → Same number of valence e⁻.

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Cations & Anions

• Cation: Positive ion (loss of electrons).
• Anion: Negative ion (gain of electrons).
• Examples:
  • Na → Na⁺ + e⁻
  • Cl + e⁻ → Cl⁻
  • Mg → Mg²⁺ + 2e⁻

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Quantum Mechanical Model

• Electrons behave as waves + particles.
• Orbitals = regions of probability for electron location.
• Example:
  • Hydrogen atom → electron in spherical 1s orbital.
  • Oxygen → electrons in 2p orbitals (dumbbell-shaped).

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Quantum Numbers

1. n (Principal) → energy level (1, 2, 3 …).
   • Ex: 2p → n=2.
2. l (Azimuthal) → subshell shape (s=0, p=1, d=2, f=3).
   • Ex: d orbital → l=2.
3. mₗ (Magnetic) → orientation (−l → +l).
   • Ex: p orbitals → −1, 0, +1.
4. mₛ (Spin) → ±½.
   • Ex: two e⁻ in same orbital = opposite spins.

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Aufbau Principle

• Electrons fill orbitals in increasing energy order.
• Sequence: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p …
• Example:
  • Oxygen (Z=8): 1s² 2s² 2p⁴.
  • Calcium (Z=20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².

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Pauli’s Exclusion Principle

• No two e⁻ can have the same 4 quantum numbers.
• Each orbital holds max 2 e⁻ with opposite spins.
• Example:
  • 1s orbital → ↑↓ (one spin +½, other spin −½).

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Hund’s Rule

• Orbitals fill singly first (with parallel spins).
• Pairing starts after all orbitals are half-filled.
• Example:
  • Carbon (Z=6): 1s² 2s² 2p² → represented as ↑ ↑ (not ↑↓).

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Atomic Radius

• Definition: Distance from nucleus → outermost electron.
• Trends:
  • Across a Period → Decreases (↑ nuclear charge pulls e⁻ closer).
  • Down a Group → Increases (new shells added).
• Examples:
  • Period: Li (152 pm) > Be (112 pm) > B (85 pm).
  • Group: Na (186 pm) < K (227 pm).

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Ionization Energy

• Definition: Energy required to remove 1 e⁻ from an atom.
• Always endothermic.
• Trends:
  • Across a Period → Increases (smaller size, stronger nucleus).
  • Down a Group → Decreases (larger size, weaker hold).
• Examples:
  • Na → 496 kJ/mol (low).
  • Cl → 1251 kJ/mol (high).

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Electron Affinity

• Definition: Energy change when an atom gains 1 e⁻.
• Usually exothermic (energy released).
• Trends:
  • Across a Period → Increases (nucleus pulls extra e⁻ strongly).
  • Down a Group → Decreases (added e⁻ farther from nucleus).
• Examples:
  • Cl: −349 kJ/mol (strong affinity).
  • Noble gases: ≈ 0 (no affinity).

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Electronegativity

• Definition: Atom’s ability to attract shared e⁻ in a bond.
• Scale: Pauling (0–4).
• Trends:
  • Across a Period → Increases (nucleus stronger).
  • Down a Group → Decreases (shielding effect).
• Examples:
  • F (4.0) → highest EN.
  • Cs (0.79) → lowest EN.

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Ionic Radius

• Cations: Smaller than neutral atom (lost shell).
• Anions: Larger than neutral atom (extra e⁻ increases repulsion).
• Trends:
  • Across a Period →
    • Cations → Decrease (nuclear charge ↑).
    • Anions → Decrease (same reason).
  • Down a Group → Increases (new shells added).
• Examples:
  • Na (186 pm) → Na⁺ (102 pm).
  • Cl (99 pm) → Cl⁻ (181 pm).

14. Elements & Electron Configuration

• Electron Configuration: Describes the distribution of electrons in an atom's orbitals.
• Example (Iron, Fe, Z=26): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ or [Ar] 4s² 3d⁶

15. Chemical Bonding (Lewis, Structure, Bond Order, EAN)

• Lewis Diagram & Structure: (See Chapter 7 above).
• Bond Order: Indicates the number of chemical bonds between a pair of atoms.
  • Formula (MOT): Bond Order = ½ (# bonding e⁻ - # antibonding e⁻)
  • Simple method: Bond Order = (# of bonds between two atoms) (e.g., O=O has a bond order of 2).
• Effective Atomic Number (EAN) Rule: EAN = Atomic No. (Z) - Oxidation State + (2 × Coordination No.)
  • Example from your PDF: For K₄[Fe(CN)₆], Fe is +2 oxidation state, CN is 6. EAN = 26 - 2 + (2 × 6) = 36 (Same as Kr, a noble gas).

Effective Atomic Number (EAN) Formula

The Effective Atomic Number (EAN) of a metal complex is calculated using the formula:

EAN = Z – X + 2Y

Where:

• Z = Atomic number of the central metal atom
• X = Oxidation number of the metal
• Y = Coordination number (number of ligands attached to the metal)

Explanation

1. Start with the total number of electrons in the neutral metal atom (Z).
2. Subtract the electrons lost due to the metal’s oxidation state (X).
3. Add the electrons donated by the ligands (2 × Y, assuming each ligand donates 2 electrons).

Note

• EAN helps predict the stability of metal complexes.
• Complexes with EAN = 18 are usually very stable, following the "18-electron rule".

Example

For a metal with:

• Z = 26 (Fe)
• X = 2 (Fe²⁺)
• Y = 6 (octahedral complex with 6 ligands)

EAN = 26 – 2 + 2×6 = 26 – 2 + 12 = 36

Of course. Based on the provided "Colloids and Colloidal Solution.pdf" (Chapter 7), here are the detailed answers to your related questions.

16. Last Chapter: All definitions, applications, classifications, physical properties of the last chapter.

The last chapter is on Colloids and Colloidal Solutions.

1. Definitions

• Colloid: A colloid is a substance microscopically dispersed throughout another substance. The particles are larger than in a solution but smaller than in a suspension, with a diameter between approximately 1 nm and 1000 nm.
• Dispersed Phase: The substance that is distributed as colloidal particles (the solute-like phase).
• Dispersion Medium: The substance in which the colloidal particles are dispersed (the solvent-like phase).

2. Classifications

Colloids are classified on several bases:

A. Based on Physical State of Dispersed Phase and Dispersion Medium:

B. Based on Nature of Interaction:

• Lyophilic Colloids (Intrinsic): The dispersed phase has a high affinity for the dispersion medium (e.g., starch, gelatin, rubber). They are self-stabilized, reversible, and heavily hydrated.
• Lyophobic Colloids (Extrinsic): The dispersed phase has no affinity for the dispersion medium (e.g., sols of gold, silver, metal hydroxides). They are unstable, irreversible, poorly hydrated, and require a stabilizing electrolyte.

C. Based on Molecular Size:

• Multimolecular Colloids: Particles are aggregates of many small atoms/molecules held by weak van der Waals forces (e.g., sulphur sol).
• Macromolecular Colloids: Particles are themselves very large molecules of colloidal size (e.g., natural polymers like proteins).
• Associated Colloids (Micelles): Substances that behave as normal electrolytes at low concentrations but form colloidal particles (micelles) at high concentrations (e.g., soaps, detergents).

3. Physical Properties

• Heterogeneity: Consist of two distinct phases.
• Filterability: Colloidal particles pass through ordinary filter paper but not through a semipermeable membrane.
• Tyndall Effect: The scattering of a beam of light by colloidal particles, making the light path visible.
• Stability: Lyophilic sols are very stable. Lyophobic sols are stable only in the absence of high electrolyte concentrations.
• Colour: Depends on particle size; larger particles absorb longer wavelengths and transmit shorter ones.

4. Electrical Properties

• Electrophoresis: The movement of charged colloidal particles towards an oppositely charged electrode under the influence of an electric field. This proves colloidal particles carry an electric charge.
• Coagulation/Flocculation: The precipitation of a colloidal solution by adding an electrolyte. The ions of the electrolyte neutralize the charge on the colloidal particles, causing them to clump together and settle.
  • Hardy-Schulze Rule: The greater the valence of the oppositely charged ion of the electrolyte, the greater its coagulating power.

5. Applications

• Medicine: Colloidal gold is used to carry drugs and antibiotics. Many medicines (like Cod Liver Oil) are emulsions.
• Food Industry: Milk, butter, cheese, ice cream, and jam are all colloidal in nature.
• Water Treatment: Sewage particles, which are colloidal, are removed using electrophoresis.
• Daily Life: Soap works by forming micelles that emulsify and trap grease. The blue color of the sky is due to the Tyndall effect caused by colloidal dust particles in the air.
• Industry: Used in paints, inks, and building roads (asphalt emulsion).

Extra Topic: How Soap Works (Micelle Formation)

This is a key example of an associated colloid.

1. Soap molecules have a hydrophobic (oil-loving) tail and a hydrophilic (water-loving) head.
2. In water, at low concentrations, they act as normal electrolytes.
3. At higher concentrations, they aggregate to form spherical clusters called micelles.
4. In a micelle, the hydrophobic tails point inwards, trapping grease and oil. The hydrophilic heads point outwards, interacting with water and keeping the micelle suspended.
5. This process emulsifies grease, allowing it to be washed away with water.