Chemistry Compendium — Quick Reference

Introduction

This final chapter is a master reference guide—a repository of essential information, formulas, mnemonics, and quick facts you'll need during exam preparation and the actual competitive exam. Think of it as your chemistry survival guide: when you're mid-exam and your mind goes blank on the activity series or pH calculations, this chapter is your anchor.

Part 1: Periodic Table Summary (Groups 1–18)

Group Characteristics

Group	Name	Valence e⁻	Common Ion	Reactivity	Key Property
1	Alkali metals	1	+1	Very high	Soft, low density
2	Alkaline earth metals	2	+2	High	Harder than Group 1
13	Boron family	3	+3	Moderate	Amphoteric oxides
14	Carbon family	4	Varies	Moderate	Tetravalency; catenation
15	Nitrogen family	5	-3, +5	Moderate	Nonmetallic character
16	Oxygen family	6	-2	High	Nonmetallic; oxidizing
17	Halogens	7	-1	Very high	F > Cl > Br > I (reactivity)
18	Noble gases	8 (or 2 for He)	None	Inert	Full valence shell

[Memory Hook] Group 1: +1 ions. Group 2: +2 ions. Group 17: -1 ions. Group 18: Inert (no ions)

Transition Metals (Groups 3–12)

Key features:

• Multiple oxidation states (e.g., Fe: +2, +3; Mn: +2, +3, +4, +7)
• Colored compounds (due to d-orbital transitions)
• Many are magnetic
• Excellent catalysts

Common transition metals:

• Fe (Iron): Oxidation states +2, +3; found in hemoglobin, steel
• Cu (Copper): Oxidation states +1, +2; red metal; excellent conductor
• Zn (Zinc): Oxidation state +2; galvanization; enzyme cofactor
• Cr (Chromium): Oxidation states +2, +3, +6; toxic +6 form
• Mn (Manganese): Oxidation states +2 to +7; permanganate (MnO₄⁻) is oxidizing agent

Part 2: Atomic Structure Quick Reference

Atomic Number, Mass Number, Electrons

Atomic Number (Z) = number of protons = number of electrons (neutral atom)

Mass Number (A) = protons + neutrons

Number of neutrons = A - Z

Electron Configuration Quick Shortcuts

Groups indicate valence electrons:

• Group 1: 1 valence electron (ns¹)
• Group 2: 2 valence electrons (ns²)
• Group 13: 3 valence electrons (ns² np¹)
• Group 14: 4 valence electrons (ns² np²)
• Group 17: 7 valence electrons (ns² np⁵)
• Group 18: 8 valence electrons (ns² np⁶, or 2 for He)

Orbital filling order (mnemonic):

Aufbau filling: s, p, d, f
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

Diagonal rule memory aid:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p
7s

[Memory Hook] Diagonal rule: Follow diagonals from top-right to bottom-left

Quantum Numbers Quick Reference

Quantum Number	Meaning	Possible Values
n	Energy level	1, 2, 3, 4, ...
l	Orbital type (s, p, d, f)	0, 1, 2, 3 (l < n)
m_l	Orbital orientation	-l to +l
m_s	Electron spin	+½ or -½

Part 3: Chemical Bonding Quick Reference

Electronegativity and Bond Type Decision Tree

EN Difference
    |
    ├─ < 0.4 → Nonpolar covalent (e.g., Cl—Cl, C—H)
    ├─ 0.4–1.7 → Polar covalent (e.g., H—Cl, H—O—H)
    └─ > 1.7 → Ionic (e.g., Na—Cl, Ca—O)

VSEPR Geometry Summary

Electron Groups	Geometry	Bond Angle	Example
2	Linear	180°	CO₂
3	Trigonal planar	120°	BF₃, NH₃ (pyramidal)
4	Tetrahedral	109.5°	CH₄, NH₃, H₂O
5	Trigonal bipyramidal	90°, 120°	PCl₅
6	Octahedral	90°	SF₆

[Memory Hook] 2 groups: linear (180°). 3 groups: trigonal (120°). 4 groups: tetrahedral (109.5°)

Hybridization Summary

Hybridization	Geometry	Bond Angle	Example
sp	Linear	180°	CO₂, BeCl₂
sp²	Trigonal planar	120°	C₂H₄, BF₃
sp³	Tetrahedral	109.5°	CH₄, NH₃
sp³d	Trigonal bipyramidal	90°, 120°	PCl₅
sp³d²	Octahedral	90°	SF₆

Part 4: Periodic Trends Quick Reference

Trends Summary

Across Period (Left → Right):
- Atomic radius: DECREASES ↓
- Ionization energy: INCREASES ↑
- Electronegativity: INCREASES ↑
- Metallic character: DECREASES ↓

Down Group (Top → Bottom):
- Atomic radius: INCREASES ↑
- Ionization energy: DECREASES ↓
- Electronegativity: DECREASES ↓
- Metallic character: INCREASES ↑

[Memory Hook] Atomic radius: smaller right-top, larger left-bottom. Reactivity (metals): higher down and left. Reactivity (nonmetals): higher right and up

Part 5: Important Chemical Formulas

Stoichiometry and Concentration

Molarity: M = moles / liters

Molality: m = moles / kg of solvent

Normality: N = equivalents / liters

Percent composition: % = (mass of element / molar mass) × 100%

Percent yield: % yield = (actual yield / theoretical yield) × 100%

Gas Laws

Ideal gas law: PV = nRT

• R = 0.0821 L·atm/(mol·K) or 8.314 J/(mol·K)

Combined gas law: P₁V₁/T₁ = P₂V₂/T₂

Dalton's law: P_total = P₁ + P₂ + ... + Pₙ

Thermodynamic Formulas

Heat: q = m × c × ΔT

• c = specific heat capacity

Enthalpy: ΔH = Σ(H products) - Σ(H reactants)

Acid-Base Chemistry

pH: pH = -log[H⁺]

pOH: pOH = -log[OH⁻]

Relationship: pH + pOH = 14 (at 25°C)

Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA])

Ka: Ka = [H⁺][A⁻] / [HA] (weak acid dissociation)

Kw: Kw = [H⁺][OH⁻] = 10⁻¹⁴ (at 25°C)

Equilibrium

Equilibrium constant: K = [products] / [reactants]

Le Chatelier's principle: System shifts to counteract disturbance

Part 6: Strong Acids and Bases (Must Memorize)

The Big 6 Strong Acids

1. HCl (Hydrochloric acid)
2. HBr (Hydrobromic acid)
3. HI (Hydroiodic acid)
4. HNO₃ (Nitric acid)
5. H₂SO₄ (Sulfuric acid)
6. HClO₄ (Perchloric acid)

Mnemonic: "The Big 6" — If you see these, assume 100% ionization

Common Strong Bases

• LiOH, NaOH, KOH (Group 1 hydroxides)
• Ca(OH)₂, Ba(OH)₂ (Some Group 2 hydroxides)

Part 7: Activity Series (Reactivity Series)

Metal Reactivity (Most to Least Reactive)

Memory aid: "Please Keep Calling My Friend Alex Zinc Can By Copper's Silver-Gold Name"

K > Na > Ca > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Hg > Ag > Pt > Au

Top (most reactive): K, Na, Ca (alkali metals, alkaline earth metals)

Middle: Transition metals (Al, Zn, Fe, Cu)

Bottom (least reactive): Ag, Au (noble metals)

Halogen Reactivity (Most to Least)

F₂ > Cl₂ > Br₂ > I₂

Fluorine is MOST reactive element overall

Part 8: Solubility Rules

General Solubility Rules

Rule	Solubility
Group 1 compounds (Li, Na, K, ...)	Soluble
NH₄⁺ compounds	Soluble
NO₃⁻, ClO₄⁻ compounds	Soluble
Cl⁻, Br⁻, I⁻ compounds	Soluble EXCEPT Ag⁺, Pb²⁺, Hg₂²⁺
SO₄²⁻ compounds	Soluble EXCEPT Ba²⁺, Pb²⁺, Ca²⁺, Sr²⁺
CO₃²⁻, PO₄³⁻, SO₃²⁻ compounds	Insoluble EXCEPT Group 1 and NH₄⁺
OH⁻ compounds	Insoluble EXCEPT Group 1, Ba(OH)₂, Ca(OH)₂ (slightly)

Quick check: If nothing says it's insoluble, it's probably soluble!

Part 9: Oxidation States (Common Examples)

Element/Ion	Oxidation State	Example
H	+1 (usually); -1 in hydrides	H₂O (+1); NaH (-1)
O	-2 (usually); -1 in peroxides	H₂O (-2); H₂O₂ (-1)
Cl	-1 (usually); +1, +3, +5, +7 in oxyacids	Cl⁻ (-1); HClO (+1); HClO₄ (+7)
Group 1	Always +1	Na⁺
Group 2	Always +2	Ca²⁺
Transition metals	Variable	Fe²⁺ (+2), Fe³⁺ (+3)

Part 10: Important Polyatomic Ions

Positive Ions (Cations)

Ion	Charge	Formula
Ammonium	+1	NH₄⁺
Hydronium	+1	H₃O⁺

Negative Ions (Anions)

Ion	Charge	Formula
Hydroxide	-1	OH⁻
Nitrate	-1	NO₃⁻
Nitrite	-1	NO₂⁻
Chlorate	-1	ClO₃⁻
Permanganate	-1	MnO₄⁻
Bicarbonate/Hydrogen carbonate	-1	HCO₃⁻
Bisulfate	-1	HSO₄⁻
Chloride	-1	Cl⁻
Sulfate	-2	SO₄²⁻
Sulfite	-2	SO₃²⁻
Carbonate	-2	CO₃²⁻
Chromate	-2	CrO₄²⁻
Dichromate	-2	Cr₂O₇²⁻
Thiosulfate	-2	S₂O₃²⁻
Phosphate	-3	PO₄³⁻
Phosphite	-3	PO₃³⁻

[Memory Hook] Common -1 anions: OH⁻, NO₃⁻, Cl⁻. Common -2: SO₄²⁻, CO₃²⁻

Part 11: Molecular Weights of Common Compounds

Compound	Molar Mass (g/mol)
H₂O	18
NaCl	58.5
CaCO₃	100
H₂SO₄	98
HCl	36.5
NH₃	17
CO₂	44
O₂	32
N₂	28
CH₄	16
C₂H₅OH (ethanol)	46
C₆H₁₂O₆ (glucose)	180

Part 12: pH of Common Substances

Substance	pH	Acidity
Battery acid	0–1	Very acidic
Lemon juice	2	Acidic
Vinegar	3	Acidic
Tomato juice	4	Acidic
Black coffee	5	Weakly acidic
Milk	6.5	Weakly acidic
Pure water	7	Neutral
Baking soda solution	8.3	Weakly basic
Ammonia solution	11	Basic
Drain cleaner	13–14	Very basic

Part 13: Acid-Base Indicators and Color Changes

Indicator	Color in Acid	Color in Base	pH Range
Methyl orange	Red	Yellow	3–4
Methyl red	Red	Yellow	5–6
Litmus	Red	Blue	5–8
Phenolphthalein	Colorless	Pink/Magenta	8–10
Methylene blue	Colorless	Blue	10–11

[Memory Hook] Phenolphthalein: colorless acid, PINK base (most important for titrations)

Part 14: Common Reactions to Remember

Combustion Reactions

General: Hydrocarbon + O₂ → CO₂ + H₂O

Examples:

• CH₄ + 2O₂ → CO₂ + 2H₂O
• C₈H₁₈ + 12.5O₂ → 8CO₂ + 9H₂O

Neutralization (Acid + Base)

H⁺ + OH⁻ → H₂O

Example: HCl + NaOH → NaCl + H₂O

Metal + Acid

General: Metal + acid → salt + H₂↑

Example: Zn + 2HCl → ZnCl₂ + H₂↑

Metal + Oxygen

General: Metal + O₂ → metal oxide

Examples:

• 4Na + O₂ → 2Na₂O (sodium burns with yellow-orange flame)
• 3Fe + 2O₂ → Fe₃O₄ (forms black powder)

Displacement Reactions

More reactive metal displaces less reactive:

Zn + CuSO₄ → ZnSO₄ + Cu (zinc displaces copper)

Cl₂ + 2KBr → 2KCl + Br₂ (chlorine displaces bromine)

Part 15: Organic Chemistry Quick Reference

Alkane Prefixes (number of carbons)

1 = meth
2 = eth
3 = prop
4 = but
5 = pent
6 = hex
7 = hept
8 = oct
9 = non
10 = dec

Functional Group Suffixes

Functional Group	Suffix	Example
Single bond	-ane	Ethane (C₂H₆)
Double bond	-ene	Ethene (C₂H₄)
Triple bond	-yne	Ethyne (C₂H₂)
Alcohol	-ol	Ethanol (C₂H₅OH)
Aldehyde	-al	Ethanal (CH₃CHO)
Ketone	-one	Propanone (CH₃COCH₃)
Carboxylic acid	-oic acid	Ethanoic acid (CH₃COOH)
Ester	-oate	Ethyl ethanoate (CH₃COOC₂H₅)
Amine	-amine	Ethanamine (C₂H₅NH₂)

Alkane Formulas

Straight-chain alkanes: CₙH₂ₙ₊₂

Alkenes (one C=C): CₙH₂ₙ

Alkynes (one C≡C): CₙH₂ₙ₋₂

Part 16: Avogadro's Number and Molar Conversions

Avogadro's Number: 6.022 × 10²³ (atoms, molecules, or formula units per mole)

Molar Mass: Mass of one mole in grams (g/mol)

Conversion triangle:

        ÷ Avogadro's number
Atoms ←────────────────────→ Moles
        × Avogadro's number
        
        ÷ molar mass (g/mol)
Mass  ←────────────────────→ Moles
        × molar mass (g/mol)

Part 17: 50 Must-Know Chemistry Facts

1. Electronegativity: F > O > N > Cl > Br > S > P > C > H > Si > Be > B > Metals
2. Strongest C-C bond: Triple bond (C≡C)
3. Strongest intermolecular force: Hydrogen bonding (H—F, H—O—, H—N—)
4. Noble gases: He, Ne, Ar, Kr, Xe, Rn (full valence shells; unreactive)
5. Halogens reactivity: F₂ > Cl₂ > Br₂ > I₂
6. Ionization energy: Increases across period; decreases down group
7. Electron affinity: Halogens highest; noble gases nearly zero
8. Atomic radius: Largest: Francium (Fr); smallest: Helium (He)
9. pH scale: 0–14; 7 is neutral; < 7 is acidic; > 7 is basic
10. Kw (water ion product): 10⁻¹⁴ at 25°C
11. Moles: 6.022 × 10²³ particles
12. STP (standard temperature, pressure): 273 K (0°C), 1 atm
13. Molar volume of ideal gas at STP: 22.4 L/mol
14. VSEPR: Electron pairs repel; arrange for maximum distance
15. sp hybridization: Linear, 180° angle
16. sp² hybridization: Trigonal planar, 120° angle
17. sp³ hybridization: Tetrahedral, 109.5° angle
18. Boyle's Law: P₁V₁ = P₂V₂ (at constant T)
19. Charles's Law: V₁/T₁ = V₂/T₂ (at constant P)
20. Gay-Lussac's Law: P₁/T₁ = P₂/T₂ (at constant V)
21. Oxidation = loss of electrons
22. Reduction = gain of electrons
23. OIL RIG: Oxidation Is Loss, Reduction Is Gain
24. Most reactive metal: Potassium (K)
25. Most reactive nonmetal: Fluorine (F)
26. Most abundant element (by mass): Oxygen (O)
27. Most abundant element (by atom count in universe): Hydrogen (H)
28. Alkali metal reactivity: Li < Na < K < Rb < Cs < Fr (increases down group)
29. Solubility: "Like dissolves like" (polar solvents dissolve polar solutes)
30. Boiling point elevation: ΔT_b = K_b × m
31. Freezing point depression: ΔT_f = K_f × m
32. Osmosis: Water moves toward higher solute concentration
33. Tyndall effect: Colloidal particles scatter light
34. Brownian motion: Random movement of colloidal particles
35. Enzyme: Biological catalyst; lowers activation energy
36. Photosynthesis: 6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂
37. Cellular respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy
38. Rusting of iron: Requires O₂, H₂O, and moisture (electrochemical process)
39. Galvanization: Coating iron with Zn (Zn is sacrificial anode)
40. Acid rain: Caused by SO₂ and NOₓ dissolving in water
41. Greenhouse effect: CO₂, CH₄, etc. trap infrared radiation
42. Ozone hole: CFCs catalytically destroy O₃ in stratosphere
43. Montreal Protocol: Phased out CFCs starting 1987
44. Green chemistry: Design processes to minimize waste and hazards
45. Polymer: Long-chain molecule of repeating monomers
46. Addition polymerization: Monomers (C=C) add without byproduct
47. Condensation polymerization: Monomers join with loss of small molecule (H₂O)
48. Plastics recycling: Symbols 1–7; 1 (PET) and 2 (HDPE) most recyclable
49. DNA: Double helix; stores genetic information; bases: A, T, G, C
50. Amino acids: Building blocks of proteins; linked by peptide bonds

Conclusion

Chemistry is vast, but these 12 chapters and this quick reference guide distill the essentials. Commit the periodic table trends, the activity series, the strong acids/bases, and the key formulas to memory. The rest will follow from understanding the fundamental principles: electron arrangement, bonding, reactivity, and equilibrium. Master these, and you'll excel not just in competitive exams but in understanding the world around you.

Good luck in your preparation. Chemistry reveals the hidden order beneath apparent chaos. That pursuit of understanding—that's what makes chemistry beautiful.

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23 MCQ Questions (Quick Reference Mastery Test)

Q1: Which of the following is NOT one of the Big 6 strong acids?

• A) HCl
• B) H₂SO₄
• C) H₃PO₄
• D) HNO₃

Q2: According to the activity series, which metal can displace copper from copper sulfate?

• A) Silver (Ag)
• B) Gold (Au)
• C) Zinc (Zn)
• D) Mercury (Hg)

Q3: Fluorine is the most reactive element. Which of the following is second most reactive?

• A) Chlorine
• B) Oxygen
• C) Nitrogen
• D) Bromine

Q4: What is the pH of a neutral solution at 25°C?

• A) 0
• B) 7
• C) 10
• D) 14

Q5: The molar volume of an ideal gas at STP (0°C, 1 atm) is:

• A) 11.2 L/mol
• B) 22.4 L/mol
• C) 44.8 L/mol
• D) 12.5 L/mol

Q6: How many electrons does a neutral sodium atom have?

• A) 10
• B) 11
• C) 12
• D) 23

Q7: According to VSEPR, what is the bond angle in a tetrahedral molecule?

• A) 90°
• B) 109.5°
• C) 120°
• D) 180°

Q8: The oxidation state of oxygen in H₂O₂ is:

• A) -2
• B) -1
• C) 0
• D) +2

Q9: Which polyatomic ion has a charge of -2?

• A) NO₃⁻
• B) OH⁻
• C) SO₄²⁻
• D) Cl⁻

Q10: In the Haber process, which compounds react?

• A) N₂ and O₂
• B) N₂ and H₂
• C) NO and H₂
• D) NH₃ and O₂

Q11: The Henderson-Hasselbalch equation calculates:

• A) pOH
• B) Kw
• C) pH of a buffer solution
• D) Osmotic pressure

Q12: Which functional group is characteristic of carboxylic acids?

• A) -OH (hydroxyl)
• B) -COOH (carboxyl)
• C) -CHO (aldehyde)
• D) -NH₂ (amino)

Q13: An alkene with 4 carbons has the molecular formula:

• A) C₄H₁₀
• B) C₄H₈
• C) C₄H₆
• D) C₄H₄

Q14: Colligative properties depend on:

• A) The identity of the solute
• B) The number of solute particles
• C) The color of the solution
• D) The solubility of the solute

Q15: Which gas is responsible for the greenhouse effect (primary)?

• A) O₂
• B) N₂
• C) CO₂
• D) Ar

Q16: In galvanization, iron is protected by:

• A) A layer of copper
• B) A layer of zinc (which oxidizes preferentially)
• C) Paint
• D) A layer of aluminum

Q17: The pH of a solution with [H⁺] = 10⁻⁴ M is:

• A) 4
• B) 10
• C) 0.0001
• D) 10⁻⁴

Q18: Phenolphthalein is colorless in acid and _____ in base.

• A) Yellow
• B) Blue
• C) Pink
• D) Red

Q19: Which polymer is formed by condensation polymerization of dicarboxylic acid and diol?

• A) Polyethylene (PE)
• B) Polyester (PET)
• C) Polystyrene (PS)
• D) Polyvinyl chloride (PVC)

Q20: The most electronegative element is:

• A) Chlorine (Cl)
• B) Oxygen (O)
• C) Fluorine (F)
• D) Nitrogen (N)

Q21: According to Avogadro's law, how many molecules are in 2 moles of O₂?

• A) 6.022 × 10²³
• B) 1.2044 × 10²⁴
• C) 3.011 × 10²³
• D) 2 × 10²³

Q22: The Big 6 strong acids include all of the following EXCEPT:

• A) HBr
• B) HI
• C) HCN
• D) HClO₄

Q23: In the Lewis structure of CO₂, carbon has how many bonding pairs of electrons?

• A) 1
• B) 2
• C) 3
• D) 4

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Answer Key: 1-C, 2-C, 3-A, 4-B, 5-B, 6-B, 7-B, 8-B, 9-C, 10-B, 11-C, 12-B, 13-B, 14-B, 15-C, 16-B, 17-A, 18-C, 19-B, 20-C, 21-B, 22-C, 23-B