Topics: Organic + Physical + Inorganic Chemistry
Total Questions: 50
Type: Single Correct MCQ
Difficulty: JEE Main Tough
Time: 60 Minutes

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SECTION A — ORGANIC CHEMISTRY

Q1.

Which compound has the highest acidic strength?

A) Phenol
B) p-Cresol
C) p-Nitrophenol
D) p-Methoxyphenol

Q2.

The correct decreasing order of electron-releasing tendency of the following groups through the +I effect is:

A) –CH₃ > –C₂H₅ > –i-C₃H₇ > –t-C₄H₉
B) –t-C₄H₉ > –i-C₃H₇ > –C₂H₅ > –CH₃
C) –C₂H₅ > –CH₃ > –t-C₄H₉ > –i-C₃H₇
D) –i-C₃H₇ > –t-C₄H₉ > –C₂H₅ > –CH₃

Q3.

Gattermann–Koch formylation of benzene requires:

A) CO + HCl / AlCl₃
B) CO + HCl / AlCl₃, CuCl
C) CO₂ + HCl / AlCl₃
D) CH₃Cl / AlCl₃

Q4.

Phenol on treatment with CHCl₃/NaOH followed by acidification gives predominantly:

A) Benzaldehyde
B) Salicylaldehyde
C) Benzoic acid
D) Salicylic acid

Q5.

Phenol treated with CO₂/NaOH under suitable conditions followed by acidification gives:

A) Benzaldehyde
B) Salicylic acid
C) Benzoic acid
D) Picric acid

Q6.

A secondary amine reacts with HNO₂ to form:

A) Alcohol + N₂
B) N-nitrosoamine
C) Diazonium salt
D) Alkene

Q7.

A tertiary aliphatic amine on treatment with HNO₂ generally forms:

A) N-nitrosoamine
B) Diazonium salt
C) Ammonium nitrite salt
D) Alcohol

Q8.

The correct order of dipole moment is:

A) HF > HCl > HBr > HI
B) HCl > HF > HBr > HI
C) HI > HBr > HCl > HF
D) HF > HBr > HCl > HI

Q9.

Which pair represents structural isomers?

A) cis-2-butene and trans-2-butene
B) n-butane and isobutane
C) D-glucose and L-glucose
D) cis-2-butene and 2-methylpropene

Q10.

Which compound can exhibit E/Z isomerism?

A) Propene
B) 2-methylpropene
C) But-2-ene
D) 2-methylbut-2-ene

Q11.

The most stable conformation of n-butane is:

A) Fully eclipsed
B) Gauche
C) Anti staggered
D) Partially eclipsed

Q12.

Ozonolysis of 2-methylpropene followed by Zn/H₂O gives:

A) Methanal + propanone
B) Methanal + ethanal
C) Ethanal + propanone
D) Two molecules of ethanal

Q13.

Which reaction proceeds predominantly through an SN2 mechanism?

A) tert-Butyl bromide + OH⁻
B) Methyl bromide + OH⁻
C) tert-Butyl chloride + H₂O
D) Benzyl alcohol + H₂O

Q14.

The rate of an SN1 reaction depends upon:

A) Substrate only
B) Nucleophile only
C) Both substrate and nucleophile
D) Solvent only

Q15.

Acid chloride reacts with water to form:

A) Ester
B) Aldehyde
C) Carboxylic acid
D) Amide

Q16.

Hydration of a terminal alkyne in the presence of HgSO₄/H₂SO₄ gives initially an enol which undergoes:

A) Aldol condensation
B) Keto-enol tautomerism
C) Cannizzaro reaction
D) Polymerisation

Q17.

Addition of HBr to propene in the presence of peroxide gives predominantly:

A) 2-Bromopropane
B) 1-Bromopropane
C) 1,2-Dibromopropane
D) Propanol

Q18.

Which reagent reduces aldehydes/ketones to hydrocarbons under strongly basic conditions?

A) NaBH₄
B) LiAlH₄
C) Zn(Hg)/HCl
D) PCC

Q19.

Wolff–Kishner reduction uses:

A) Zn(Hg)/HCl
B) NH₂NH₂/KOH, heat
C) NaBH₄
D) H₂/Ni

Q20.

A nitrile reacts with excess Grignard reagent followed by hydrolysis to give:

A) Primary alcohol
B) Secondary alcohol
C) Ketone
D) Tertiary amine

Q21.

Baeyer’s reagent is:

A) Acidified KMnO₄
B) Cold dilute alkaline KMnO₄
C) Br₂/CCl₄
D) Ammoniacal AgNO₃

Q22.

Which compound gives a positive ammoniacal AgNO₃ test?

A) Ethene
B) Ethyne
C) Propene
D) Benzene

Q23.

Cold dilute KMnO₄ reacts with an alkene mainly to produce:

A) Alkane
B) Vicinal diol
C) Aldehyde only
D) Ketone only

Q24.

Which compound can undergo Cannizzaro reaction?

A) CH₃CHO
B) C₂H₅CHO
C) HCHO
D) CH₃COCH₃

Q25.

Which pair can undergo aldol condensation?

A) Benzaldehyde + benzaldehyde
B) Formaldehyde + formaldehyde
C) Acetaldehyde + acetaldehyde
D) Benzaldehyde + formaldehyde

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SECTION B — PHYSICAL CHEMISTRY

Q26.

The equivalent weight of H₂SO₄ in an acid-base reaction involving complete neutralisation is:

A) 98
B) 49
C) 24.5
D) 196

Q27.

A solution contains 0.2 equivalent of solute in 500 mL solution. Its normality is:

A) 0.1 N
B) 0.2 N
C) 0.4 N
D) 2.5 N

Q28.

For a hydrogen-like species, the energy of an electron in the nth orbit is proportional to:

A) n²
B) 1/n
C) –1/n²
D) –n²

Q29.

For a hydrogen atom, the energy gap between n = 1 and n = 2 is:

A) Greater than that between n = 2 and n = 3
B) Equal to that between n = 2 and n = 3
C) Less than that between n = 2 and n = 3
D) Zero

Q30.

For a first-order reaction, the time required for 100% completion is:

A) t₁/₂
B) 2t₁/₂
C) Finite
D) Infinite

Q31.

For a zero-order reaction, if initial concentration is [A]₀, the time for complete consumption is:

A) [A]₀/k
B) k/[A]₀
C) 0.693/k
D) ∞

Q32.

For an ideal gas undergoing reversible isothermal expansion:

A) w = 0
B) Work is maximum
C) Work is minimum
D) ΔU > 0

Q33.

Compared with irreversible expansion against constant external pressure between the same initial and final states, reversible expansion gives:

A) Less work by gas
B) More work by gas
C) Same work
D) Zero work

Q34.

At constant pressure, heat supplied to a system is equal to:

A) ΔU
B) ΔH
C) ΔS
D) ΔG

Q35.

At constant volume, heat supplied to a system is equal to:

A) ΔH
B) ΔU
C) ΔG
D) ΔS

Q36.

For a reaction, ΔH = –100 kJ mol⁻¹ and Δn(g) = –2. At 300 K, approximately:$$\Delta H-\Delta U = \Delta n_gRT$$

Therefore ΔU is:

A) –95.0 kJ mol⁻¹
B) –100 kJ mol⁻¹
C) –105.0 kJ mol⁻¹
D) +105 kJ mol⁻¹

Q37.

The heat capacity of a body is 400 J K⁻¹ and its mass is 200 g. Its specific heat capacity is:

A) 0.5 J g⁻¹ K⁻¹
B) 2 J g⁻¹ K⁻¹
C) 200 J g⁻¹ K⁻¹
D) 800 J g⁻¹ K⁻¹

Q38.

For an exothermic reaction:

A) ΔH > 0
B) ΔH < 0
C) ΔH = 0 always
D) ΔS must be negative

Q39.

Hess’s law is based on the fact that enthalpy is a:

A) Path function
B) State function
C) Intensive property only
D) Rate-dependent property

Q40.

For a spontaneous process at constant temperature and pressure:

A) ΔG > 0
B) ΔG = 0 always
C) ΔG < 0
D) ΔH must be positive

Q41.

Entropy generally increases when:

A) Gas → liquid
B) Liquid → solid
C) Solid → gas
D) Gas → solid

Q42.

If Ksp of AgCl is 1.0 × 10⁻¹⁰, its solubility in pure water is approximately:

A) 10⁻⁵ M
B) 10⁻¹⁰ M
C) 10⁻² M
D) 10⁻¹⁵ M

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SECTION C — INORGANIC CHEMISTRY

Q43.

The correct spectrochemical series is:

A) I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < CN⁻
B) CN⁻ < NH₃ < H₂O < F⁻ < Cl⁻ < Br⁻ < I⁻
C) F⁻ < I⁻ < Br⁻ < Cl⁻ < H₂O < CN⁻ < NH₃
D) NH₃ < H₂O < F⁻ < Cl⁻ < Br⁻ < I⁻ < CN⁻

Q44.

The molecule CO according to MOT has:

A) Bond order 1
B) Bond order 2
C) Bond order 3
D) Bond order 0

Q45.

Which species is paramagnetic?

A) N₂
B) O₂
C) F₂
D) Ne₂

Q46.

The oxidation state of Mn in KMnO₄ is:

A) +2
B) +4
C) +6
D) +7

Q47.

A reaction in which the same element is simultaneously oxidised and reduced is called:

A) Displacement
B) Comproportionation
C) Disproportionation
D) Neutralisation

Q48.

In a comproportionation reaction:

A) Same oxidation state gives two different oxidation states
B) Two different oxidation states combine to give an intermediate oxidation state
C) Only oxidation occurs
D) Only reduction occurs

Q49.

During electrolysis of molten NaCl, the product at the cathode is:

A) Cl₂
B) NaOH
C) Na
D) H₂

Q50.

During electrolysis of aqueous CuSO₄ using inert electrodes, the cathode reaction is predominantly:

A) 2H₂O → O₂ + 4H⁺ + 4e⁻
B) Cu²⁺ + 2e⁻ → Cu
C) SO₄²⁻ → SO₂
D) H₂O → H₂

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ANSWER KEY

Q	Ans	Q	Ans	Q	Ans	Q	Ans
1	C	14	A	27	C	40	C
2	B	15	C	28	C	41	C
3	B	16	B	29	A	42	A
4	B	17	B	30	D	43	A
5	B	18	C	31	A	44	C
6	B	19	B	32	B	45	B
7	C	20	C	33	B	46	D
8	A	21	B	34	B	47	C
9	B	22	B	35	B	48	B
10	C	23	B	36	C	49	C
11	C	24	C	37	B	50	B
12	A	25	C	38	B
13	B	26	B	39	B

Carbon and Hydrogen

Carius method

Dumas’ Method for Estimation of Nitrogen

Percentageof sulphur

m1= mass of BaSO₄

m= mass of organic compound

Percentage of Phosphorus

Percentage of Oxygen

Chapters: Solutions • Electrochemistry • Chemical Kinetics

Maximum Marks: 30
Time: 1 Hour

General Instructions

1. All questions are numerical/problem-based.
2. Show the complete calculation and appropriate units.
3. Use log 2 = 0.3010, log 3 = 0.4771 wherever required.
4. Take F = 96500 C mol⁻¹ and R = 8.314 J K⁻¹ mol⁻¹, wherever required.

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SECTION A — 1 × 6 = 6 Marks

Answer all questions.

Q1. Calculate the molality of a solution prepared by dissolving 18 g glucose (M = 180 g mol⁻¹) in 180 g of water.

Q2. The vapour pressure of pure water at 300 K is 31.8 mm Hg. Calculate the vapour pressure of a solution containing 1 mole of glucose in 9 moles of water.

Q3. Calculate the EMF of the cell at 298 K:

Zn | Zn²⁺ (0.1 M) || Cu²⁺ (0.01 M) | Cu

Given: E°cell = 1.10 V.

Q4. A current of 2 A is passed through a CuSO₄ solution for 30 minutes. Calculate the mass of copper deposited.
(Molar mass of Cu = 63.5 g mol⁻¹)

Q5. A first-order reaction has a half-life of 20 minutes. Calculate its rate constant.

Q6. For a reaction, the rate constant is 2.0 × 10⁻³ s⁻¹. Calculate its half-life.

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SECTION B — 2 × 6 = 12 Marks

Answer all questions.

Q7. 5.85 g NaCl is dissolved in 500 g water. Calculate the depression in freezing point.
Given: Kf = 1.86 K kg mol⁻¹.
Assume complete dissociation of NaCl.

Q8. Calculate the molarity of a solution prepared by dissolving 4.9 g H₂SO₄ in water to make the final volume 250 mL.
(Molar mass of H₂SO₄ = 98 g mol⁻¹)

Q9. Calculate the standard Gibbs energy change for a cell reaction involving n = 2 electrons, if E°cell = 1.10 V.
Given: F = 96500 C mol⁻¹.

Q10. The conductivity of a 0.02 M KCl solution is 2.5 × 10⁻³ S cm⁻¹. Calculate its molar conductivity.

Q11. For a first-order reaction, 75% of the reactant decomposes in 40 minutes. Calculate the rate constant.

Q12. A reaction follows second-order kinetics. Its rate constant is 0.5 L mol⁻¹ s⁻¹ and initial concentration is 0.2 mol L⁻¹. Calculate the half-life of the reaction.

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SECTION C — 3 × 4 = 12 Marks

Answer all questions.

Q13. A solution is prepared by dissolving 10 g of a non-volatile solute in 90 g of water. The vapour pressure of pure water at a certain temperature is 40 mm Hg, while that of the solution is 38 mm Hg.

Calculate the molar mass of the solute.

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Q14. A conductivity cell containing 0.01 M KCl solution has a resistance of 100 Ω. The conductivity of the solution is 1.41 × 10⁻³ S cm⁻¹.

Calculate:

a) Cell constant
b) Conductance of the solution
c) Molar conductivity of KCl

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Q15. A galvanic cell is represented as:

Mg | Mg²⁺ (0.01 M) || Cu²⁺ (0.001 M) | Cu

Given:

E°(Mg²⁺/Mg) = −2.37 V
E°(Cu²⁺/Cu) = +0.34 V

Calculate:

a) E°cell
b) Ecell at 298 K
c) ΔG° for the cell reaction

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Q16. The decomposition of a substance follows first-order kinetics. Its concentration decreases from 0.80 M to 0.20 M in 30 minutes.

Calculate:

a) Rate constant
b) Half-life
c) Time required for the concentration to decrease from 0.20 M to 0.05 M

When a ray of light travels from one transparent medium to another, its speed changes, causing the ray to change its direction. This phenomenon is called refraction of light.

Consider a light ray travelling through air → water → glass → air.

1. Air to Water

When light enters water from air, it travels from a rarer medium to a denser medium. Therefore, the ray bends towards the normal.

2. Water to Glass

When the ray travels from water to glass, it again enters a denser medium because the refractive index of glass is greater than that of water. Hence, the ray bends towards the normal again.

3. Glass to Air

Finally, when the ray emerges from glass into air, it travels from a denser medium to a rarer medium. Therefore, it bends away from the normal.

Important Result

If the first and last surfaces are parallel, the final emergent ray is parallel to the original incident ray, although it is shifted sideways. This sideways shift is called lateral displacement.

Remember for JEE/NEET:

Rarer → Denser → Towards normal
Denser → Rarer → Away from normal

The amount of bending depends on the refractive indices of the media and the angle at which light strikes the surface.

P-BLOCK GROUP 13 & 14 — 50 MCQs

JEE + NEET Practice

GROUP 13 — BORON FAMILY

1. The general electronic configuration of Group 13 elements is:

A) $ns^2np^1$
B) $ns^2np^2$
C) $ns^1np^2$
D) $ns^2np^3$

2. Which element of Group 13 shows the maximum covalent character?

A) B
B) Al
C) Ga
D) Tl

3. The most important reason for the anomalous behaviour of boron is:

A) Large size
B) Low ionisation enthalpy
C) Small size and high ionisation enthalpy
D) Metallic character

4. Which Group 13 element is a metalloid?

A) B
B) Al
C) Ga
D) Tl

5. The correct order of atomic radii in Group 13 is:

A) B < Al < Ga < In < Tl
B) B < Ga < Al < In < Tl
C) B < Al < In < Ga < Tl
D) Al < B < Ga < In < Tl

6. Which element exhibits the maximum tendency for the +1 oxidation state?

A) B
B) Al
C) Ga
D) Tl

7. The inert pair effect is maximum in:

A) B
B) Al
C) Ga
D) Tl

8. Which of the following is electron deficient?

A) $NH_3$
B) $BF_3$
C) $CH_4$
D) $H_2O$

9. The hybridisation of boron in $BF_3$ is:

A) $sp$
B) $sp^2$
C) $sp^3$
D) $dsp^2$

10. $BF_3$ acts as:

A) Lewis acid
B) Lewis base
C) Brønsted acid only
D) Oxidising agent only

11. The shape of $BF_3$ is:

A) Linear
B) Trigonal planar
C) Tetrahedral
D) Pyramidal

12. Which is the strongest Lewis acid?

A) $BF_3$
B) $BCl_3$
C) $BBr_3$
D) $BI_3$

13. The Lewis acidity of boron trihalides follows:

A) $BF_3>BCl_3>BBr_3>BI_3$
B) $BF_3<BCl_3<BBr_3<BI_3$
C) $BF_3=BCl_3=BBr_3=BI_3$
D) $BCl_3>BF_3>BBr_3>BI_3$

14. The unusual Lewis acidity order of boron halides is mainly due to:

A) Electronegativity only
B) Back bonding
C) Hydrogen bonding
D) Steric effect only

15. Borax has the formula:

A) $NaBO_2$
B) $Na_2B_4O_7\cdot10H_2O$
C) $H_3BO_3$
D) $Na_3BO_3$

16. Boric acid is:

A) Tribasic
B) Dibasic
C) Monobasic
D) Strong acid

17. Boric acid behaves as a Lewis acid because it:

A) Donates $H^+$ directly
B) Accepts $OH^-$
C) Donates $OH^-$
D) Accepts $H^+$

18. On heating boric acid, the final product formed is:

A) $B_2O_3$
B) $B_2O_5$
C) $BO_2$
D) $B_2O$

19. Aluminium reacts with NaOH to form:

A) $AlCl_3$
B) Sodium aluminate
C) $Al_2O_3$ only
D) $Al(OH)_3$ only

20. Which oxide is amphoteric?

A) $B_2O_3$
B) $Al_2O_3$
C) $CO_2$
D) $SO_2$

21. Aluminium becomes passive in concentrated $HNO_3$ because:

A) Al dissolves rapidly
B) A protective oxide layer forms
C) Aluminium nitrate evaporates
D) Hydrogen is evolved

22. Cryolite used in the extraction of aluminium is:

A) $AlF_3$
B) $Na_3AlF_6$
C) $NaAlO_2$
D) $Al_2O_3$

23. The main purpose of cryolite in the Hall-Héroult process is to:

A) Increase melting point
B) Lower melting point and increase conductivity
C) Act as reducing agent
D) Produce aluminium chloride

24. Which of the following is NOT a property of boron?

A) It is a metalloid
B) It forms electron-deficient compounds
C) It readily forms $B^{3+}$ ions in aqueous solution
D) It forms covalent compounds

25. Diborane has the molecular formula:

A) $BH_3$
B) $B_2H_6$
C) $B_2H_4$
D) $BH_4$

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GROUP 14 — CARBON FAMILY

26. General electronic configuration of Group 14 elements is:

A) $ns^2np^1$
B) $ns^2np^2$
C) $ns^2np^3$
D) $ns^1np^3$

27. Which element shows the maximum catenation?

A) C
B) Si
C) Ge
D) Sn

28. The exceptional catenation ability of carbon is mainly due to:

A) Large size
B) Strong C–C bond
C) Metallic character
D) Low ionisation enthalpy

29. The most stable oxidation state of carbon is:

A) +4 only
B) +2 only
C) −4 only
D) Depends on compound

30. Down Group 14, the stability of +2 oxidation state:

A) Decreases
B) Increases
C) Remains constant
D) First increases then decreases

31. The inert pair effect in Group 14 is maximum for:

A) C
B) Si
C) Sn
D) Pb

32. Which compound is a greenhouse gas?

A) CO₂
B) CO
C) SiC
D) $CH_4$

33. Carbon monoxide is:

A) Strongly acidic oxide
B) Neutral oxide
C) Strongly basic oxide
D) Amphoteric oxide

34. Carbon dioxide is:

A) Basic oxide
B) Acidic oxide
C) Neutral oxide
D) Amphoteric oxide

35. Which of the following is an allotrope of carbon?

A) Quartz
B) Graphite
C) Silica
D) Silicate

36. In diamond, each carbon atom is:

A) $sp$-hybridised
B) $sp^2$-hybridised
C) $sp^3$-hybridised
D) $dsp^2$-hybridised

37. In graphite, each carbon atom is:

A) $sp$-hybridised
B) $sp^2$-hybridised
C) $sp^3$-hybridised
D) $sp^3d$-hybridised

38. Graphite conducts electricity because:

A) It contains ionic bonds
B) It has delocalised electrons
C) Carbon is metallic
D) It contains free protons

39. Which allotrope of carbon is the hardest naturally occurring substance?

A) Graphite
B) Diamond
C) Fullerene
D) Graphene

40. Silicon dioxide has:

A) Linear molecular structure
B) Tetrahedral 3-D network structure
C) Ionic structure
D) Planar molecular structure

41. The basic structural unit of $SiO_2$ is:

A) $SiO_6$ octahedron
B) $SiO_4$ tetrahedron
C) $SiO_3$ triangle
D) $SiO_2$ molecule

42. Which oxide is acidic?

A) $CO$
B) $CO_2$
C) $SnO$
D) PbO

43. Which Group 14 element has the strongest tendency to form $+2$ oxidation state?

A) C
B) Si
C) Sn
D) Pb

44. The reaction of $CO_2$ with lime water initially gives:

A) $CaCO_3$
B) $Ca(HCO_3)_2$
C) $CaO$
D) $Ca(OH)_2$

45. On passing excess $CO_2$ through lime water, the solution becomes clear because:

A) $CaCO_3$ converts into soluble $Ca(HCO_3)_2$
B) $CaCO_3$ decomposes to CaO
C) $CO_2$ escapes
D) CaO precipitates

46. Which of the following is used as a semiconductor?

A) Diamond
B) Silicon
C) Graphite
D) CO₂

47. Silicon reacts with nitrogen at high temperature to form:

A) $SiN$
B) $Si_3N_4$
C) $Si_2N_3$
D) $SiN_2$

48. Which of the following is a carbide?

A) $SiC$
B) $SiO_2$
C) $SiCl_4$
D) $SiH_4$

49. Silicones are polymers containing repeating:

A) C–C–C chain only
B) Si–O–Si linkage
C) Si–Si–Si linkage only
D) C–O–C linkage only

50. Which statement is correct?

A) CO₂ is linear and non-polar
B) CO₂ is bent and polar
C) SiO₂ is a discrete linear molecule
D) Diamond conducts electricity efficiently

Answer Key

Q	Ans	Q	Ans	Q	Ans	Q	Ans	Q	Ans
1	A	11	B	21	B	31	D	41	B
2	A	12	A	22	B	32	A	42	B
3	C	13	C	23	B	33	B	43	D
4	A	14	B	24	C	34	B	44	A
5	B	15	B	25	B	35	B	45	A
6	D	16	C	26	B	36	C	46	B
7	D	17	B	27	A	37	B	47	B
8	B	18	A	28	B	38	B	48	A
9	B	19	B	29	D	39	B	49	B
10	A	20	B	30	B	40	B	50	A

A. Molten salts

1. Molten NaCl

$$\mathrm{NaCl(l)\rightarrow Na^+ + Cl^-}$$

Cathode:$$\mathrm{Na^+ +e^-\rightarrow Na}$$

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Answer: A) Na, Cl₂

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2. Molten CaCl₂

Cathode:$$\mathrm{Ca^{2+}+2e^-\rightarrow Ca}$$

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Answer: A) Ca + Cl₂

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3. Molten Al₂O₃

At cathode:$$\mathrm{Al^{3+}+3e^-\rightarrow Al}$$

At anode:$$\mathrm{2O^{2-}\rightarrow O_2+4e^-}$$

So theoretically:$$\boxed{\mathrm{Al+O_2}}$$

Answer: A

Important industrial note: In the Hall-Héroult process, carbon anodes are used, so oxygen reacts with carbon and mainly forms CO₂ (with some CO).

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4. Why molten NaCl instead of aqueous NaCl for sodium?

In aqueous NaCl, both Na⁺ and water can be reduced at the cathode.

Water is preferentially reduced:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

rather than:$$\mathrm{Na^++e^-\rightarrow Na}$$

Therefore sodium metal cannot normally be obtained from aqueous NaCl.

In molten NaCl, water is absent, so:$$\boxed{\mathrm{Na^+\rightarrow Na}}$$

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B. Aqueous NaCl / Brine

5. Electrolysis of concentrated aqueous NaCl

Cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Overall:$$\boxed{\mathrm{2NaCl+2H_2O\rightarrow 2NaOH+H_2+Cl_2}}$$

Answer: B) H₂ at cathode, Cl₂ at anode

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6. Solution left after electrolysis of brine

Na⁺ remains in solution and combines with OH⁻:$$\mathrm{Na^++OH^-\rightarrow NaOH}$$

Therefore:$$\boxed{\mathrm{NaOH}}$$

Answer: A

This is the basis of the chlor-alkali process.

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7. Why isn’t Na⁺ discharged in aqueous NaCl?

Because reduction of water is thermodynamically/kinetically more favorable than reduction of Na⁺ under ordinary aqueous electrolysis conditions.$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Hence:$$\boxed{\mathrm{H_2\ rather\ than\ Na}}$$

---

8. Dilute NaCl solution

At cathode:$$\boxed{H_2}$$

At anode, especially in dilute solution:$$\boxed{O_2}$$

because oxidation of water/OH⁻ competes strongly with chloride oxidation.

JEE caution: The anodic product from chloride solutions depends on concentration and electrode conditions. Concentrated brine favors Cl₂.

---

9. pH near cathode in NaCl electrolysis

Cathode reaction:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

OH⁻ concentration increases.

Therefore:$$\boxed{\text{pH increases; solution becomes basic}}$$

---

C. Aqueous CuSO₄

10. CuSO₄ with Cu electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

Thus:$$\boxed{\text{Cu deposits at cathode and Cu dissolves at anode}}$$

Answer: A

---

11. CuSO₄ with Pt electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Therefore:$$\boxed{\mathrm{Cu\ at\ cathode,\ O_2\ at\ anode}}$$

Answer: A

---

12. What happens to CuSO₄ concentration with Pt electrodes?

Cu²⁺ is removed:$$\mathrm{Cu^{2+}\rightarrow Cu}$$

while sulfate remains in solution.

Therefore CuSO₄ concentration decreases.

Also H₂SO₄ effectively accumulates as electrolysis proceeds.$$\boxed{[\mathrm{CuSO_4}]\downarrow}$$

---

13. CuSO₄ with Cu electrodes

Cathode removes Cu²⁺:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode produces Cu²⁺:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

The rates are equal ideally.

Therefore:$$\boxed{\text{CuSO₄ concentration remains approximately unchanged}}$$

---

14. Which electrode loses mass?

At anode:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

So copper anode dissolves.$$\boxed{\text{Anode}}$$

---

D. Aqueous AgNO₃

15. AgNO₃ with Pt electrodes

Cathode:$$\mathrm{Ag^++e^-\rightarrow Ag}$$

Anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Therefore:$$\boxed{\mathrm{Ag\ at\ cathode,\ O_2\ at\ anode}}$$

Answer: A

---

16. AgNO₃ with Ag electrodes

Cathode:$$\mathrm{Ag^++e^-\rightarrow Ag}$$

Anode:$$\mathrm{Ag\rightarrow Ag^++e^-}$$

Thus:$$\boxed{\text{Ag deposits at cathode and Ag dissolves at anode}}$$

Answer: A

---

17. Why is Ag⁺ discharged instead of water?

Reduction of Ag⁺:$$\mathrm{Ag^++e^-\rightarrow Ag}$$

has a much more favorable reduction potential than reduction of water under ordinary conditions.

Therefore:$$\boxed{\mathrm{Ag^+\ is\ preferentially\ reduced}}$$

---

E. Aqueous KI / KBr / NaCl

18. Aqueous KI

Cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Anode:$$\mathrm{2I^-\rightarrow I_2+2e^-}$$

Therefore:$$\boxed{\mathrm{H_2+I_2}}$$

Answer: B

---

19. Species discharged at anode in KI

I⁻ is easily oxidized:$$\mathrm{2I^-\rightarrow I_2+2e^-}$$

Answer: B) I⁻

---

20. Increasing ease of oxidation

Halide oxidation tendency:$$\boxed{\mathrm{F^-<Cl^-<Br^-<I^-}}$$

Thus I⁻ is oxidized most easily.

---

21. Aqueous NaBr

Cathode:$$\mathrm{H_2O\rightarrow H_2}$$

Anode:$$\mathrm{2Br^-\rightarrow Br_2+2e^-}$$

Therefore:$$\boxed{\mathrm{H_2+Br_2}}$$

Answer: B

---

22. Why concentrated NaCl gives Cl₂ but dilute solution may give O₂?

At the anode, two processes compete:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

and$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

High Cl⁻ concentration favors chloride oxidation, while in dilute chloride solution oxygen evolution becomes more competitive.$$\boxed{\text{Concentrated Cl⁻ → Cl₂ favored}}$$$$\boxed{\text{Dilute Cl⁻ → O₂ can become favored}}$$

---

F. Selective discharge

23. Which ion is preferentially discharged?

Among:$$\mathrm{Na^+,Cu^{2+},Ag^+,H^+}$$

Ag⁺ is easiest to reduce.$$\boxed{\mathrm{Ag^++e^-\rightarrow Ag}}$$

Answer: Ag⁺

---

24. Increasing tendency to get reduced

Using standard reduction potentials:$$\boxed{\mathrm{Na^+<Mg^{2+}<Zn^{2+}<Cu^{2+}<Ag^+}}$$

So Ag⁺ has the greatest tendency to be reduced.

---

25. Most easily oxidized

$$\mathrm{F^-,Cl^-,Br^-,I^-}$$

Oxidation tendency:$$\boxed{\mathrm{I^->Br^->Cl^->F^-}}$$

Answer: I⁻

---

26. Why isn’t F⁻ discharged in aqueous solution?

Fluoride is extremely difficult to oxidize.

Water oxidation to O₂ occurs preferentially under ordinary aqueous electrolysis conditions.$$\boxed{\mathrm{F^-\ does\ not\ normally\ give\ F_2}}$$

---

27. Aqueous Na₂SO₄ with inert electrodes

Na⁺ is not reduced.

Water is reduced:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

At anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Therefore:$$\boxed{H_2+O_2}$$

Answer: B

---

28. Aqueous KOH

At cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

At anode:$$\mathrm{4OH^-\rightarrow O_2+2H_2O+4e^-}$$

Therefore:$$\boxed{H_2+O_2}$$

Answer: B

---

29. Dilute H₂SO₄

Cathode:$$\boxed{H_2}$$

Anode:$$\boxed{O_2}$$

Answer: B

Sulphate ion generally remains a spectator.

---

G. Mixed electrolytes

30. Cu²⁺ and Ag⁺ together

Ag⁺ has greater reduction tendency:$$\mathrm{Ag^++e^-\rightarrow Ag}$$

occurs before Cu²⁺ reduction.$$\boxed{\mathrm{Ag^+\ first}}$$

---

31. Ag⁺, Cu²⁺, H⁺ and Na⁺

General reduction preference:$$\boxed{\mathrm{Ag^+>Cu^{2+}>H^+>Na^+}}$$

Thus:

1. Ag⁺ → Ag
2. Cu²⁺ → Cu
3. H⁺/water → H₂
4. Na⁺ generally remains in solution

---

32. Cl⁻, Br⁻ and I⁻

Oxidation preference:$$\boxed{\mathrm{I^->Br^->Cl^-}}$$

Thus I₂ forms first, then Br₂, then Cl₂ as conditions change.

---

33. Cu²⁺, H⁺ and Na⁺

Cu²⁺ is preferentially reduced:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

After Cu²⁺ becomes sufficiently depleted, H₂ evolution may occur.$$\boxed{\mathrm{Cu^{2+}\ first}}$$

---

34. Cu²⁺ and Zn²⁺

Cu²⁺ has much higher reduction potential:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Therefore:$$\boxed{\mathrm{Cu}}$$

is deposited preferentially.

---

35. When Cu²⁺ concentration becomes very low

The effective reduction potential for Cu²⁺ decreases according to the Nernst equation.

Eventually hydrogen evolution may become competitive:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Thus:$$\boxed{\mathrm{H_2\ may\ start\ evolving}}$$

---

H. Electrolysis of water

36. Volume ratio of H₂ : O₂

Overall:$$\mathrm{2H_2O\rightarrow2H_2+O_2}$$

Therefore:$$\boxed{H_2:O_2=2:1}$$

Answer: B

---

37. Gas at cathode

Reduction occurs at cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$$$\boxed{H_2}$$

---

38. Gas at anode

Oxidation occurs:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$$$\boxed{O_2}$$

---

39. Why add electrolyte to water?

Pure water has very low electrical conductivity.

Adding a suitable electrolyte increases the number of ions and therefore increases conductivity.$$\boxed{\text{Electrolyte increases conductivity}}$$

---

I. Quantitative JEE questions

Use:$$\boxed{m=\frac{MIt}{nF}}$$

where

• $M$ = molar mass
• $I$ = current
• $t$ = time
• $n$ = electrons
• $F=96500\ C\,mol^{-1}$

---

40. 2 A through CuSO₄ for 965 s

For Cu:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$$$m=\frac{63.5\times2\times965}{2\times96500}$$$$\boxed{m=0.635\ g}$$

---

41. 1.93 A through AgNO₃ for 1000 s

$$\mathrm{Ag^++e^-\rightarrow Ag}$$$$m=\frac{108\times1.93\times1000}{96500}$$$$\boxed{m=2.16\ g}$$

---

42. Faradays required for 1 mol Al

$$\mathrm{Al^{3+}+3e^-\rightarrow Al}$$

1 mol Al requires 3 mol electrons.$$\boxed{3F}$$

---

43. Faradays required for 1 mol H₂

$$\mathrm{2H^++2e^-\rightarrow H_2}$$

2 mol electrons are needed.$$\boxed{2F}$$

---

44. Faradays required for 1 mol O₂

Oxidation:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Therefore:$$\boxed{4F}$$

---

45. 9650 C through acidified water

$$1F=96500C$$

Therefore:$$9650C=0.1F$$

For H₂:$$2F\rightarrow1\ mol\ H_2$$

So:$$0.1F\rightarrow0.05\ mol\ H_2$$

At STP:$$V=0.05\times22.4$$$$\boxed{V=1.12L}$$

---

46. Ag and Cu cells in series

Given:$$m_{\mathrm{Ag}}=1.08g$$

Moles Ag:$$n=\frac{1.08}{108}=0.01$$

Ag requires 1 electron per ion:$$0.01\ mol\ e^-$$

For Cu²⁺:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Therefore Cu deposited:$$0.005\ mol$$

Mass:$$0.005\times63.5$$$$\boxed{0.3175g}$$

---

47. Equal electricity through molten NaCl and AlCl₃

Na⁺ needs 1 electron:$$\mathrm{Na^++e^-\rightarrow Na}$$

Al³⁺ needs 3:$$\mathrm{Al^{3+}+3e^-\rightarrow Al}$$

Equivalent weights:$$E_{\mathrm{Na}}=23$$$$E_{\mathrm{Al}}=\frac{27}{3}=9$$

Therefore:$$\boxed{m_{\mathrm{Na}}:m_{\mathrm{Al}}=23:9}$$

---

48. 5 A for 20 minutes through CuSO₄

$$t=20\times60=1200s$$$$m=\frac{63.5\times5\times1200}{2\times96500}$$$$\boxed{m\approx1.975g}$$

---

J. Electrolysis in series

49. 0.108 g Ag deposited

Moles Ag:$$\frac{0.108}{108}=0.001$$

Thus:$$0.001\ mol\ e^-$$

For Cu:$$n_{\mathrm{Cu}}=\frac{0.001}{2}=0.0005$$

Mass:$$0.0005\times63.5$$$$\boxed{0.03175g}$$

---

50. AgNO₃, CuSO₄ and AlCl₃ in series

For equal electricity:$$m\propto\frac{M}{n}$$

Ag:$$\frac{108}{1}=108$$

Cu:$$\frac{63.5}{2}=31.75$$

Al:$$\frac{27}{3}=9$$

Therefore:$$\boxed{\mathrm{Ag:Cu:Al=108:31.75:9}}$$

---

51. Which gives maximum mass?

For equal electricity:$$m\propto\frac{M}{n}$$

Ag has the largest equivalent weight.$$\boxed{\mathrm{AgNO_3}}$$

---

52. Equal electricity through AgNO₃, CuSO₄ and Al₂(SO₄)₃

Equivalent weights:$$Ag=108$$$$Cu=\frac{63.5}{2}=31.75$$$$Al=\frac{27}{3}=9$$

Therefore:$$\boxed{Ag>Cu>Al}$$

---

K. Faraday + products

53. 0.1 mol Cu deposited from CuSO₄

Reaction:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

0.1 mol Cu requires:$$0.2\ mol\ e^-$$

At anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

4 electrons → 1 mol O₂.

Therefore:$$0.2e^-\rightarrow0.05\ mol\ O_2$$$$\boxed{0.05\ mol\ O_2}$$

---

54. 2 mol H₂ produced

$$\mathrm{2H_2O\rightarrow2H_2+O_2}$$

Ratio:$$H_2:O_2=2:1$$

Therefore:$$\boxed{1\ mol\ O_2}$$

---

55. 1 mol Al produced

$$\mathrm{Al^{3+}+3e^-\rightarrow Al}$$

1 mol Al requires 3 mol electrons.

For O₂:$$4e^-\rightarrow1mol O_2$$

Therefore:$$n(O_2)=\frac34$$$$\boxed{0.75\ mol\ O_2}$$

---

56. 2 mol Na produced

$$\mathrm{Na^++e^-\rightarrow Na}$$

2 mol Na requires 2 mol electrons.

At anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Thus:$$\boxed{1\ mol\ Cl_2}$$

---

57. 2 mol H₂ produced in NaCl electrolysis

Cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

2 mol H₂ requires 4 mol electrons.

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

4 electrons produce:$$\boxed{2\ mol\ Cl_2}$$

---

L. Electrolysis + pH

58. Cathode region in aqueous NaCl

OH⁻ is produced:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Therefore:$$\boxed{\text{Basic}}$$

Answer: B

---

59. Electrolysis of dilute H₂SO₄

Overall:$$\mathrm{2H_2O\rightarrow2H_2+O_2}$$

H₂SO₄ mainly acts as an electrolyte and is not consumed overall.

Therefore its concentration can remain approximately constant if water loss/volume effects are neglected.$$\boxed{\text{pH does not change significantly ideally}}$$

---

60. CuSO₄ + Pt electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

H⁺ accumulates.

Therefore:$$\boxed{\text{acidity increases}}$$

---

61. AgNO₃ + Pt electrodes

Cathode:$$\mathrm{Ag^++e^-\rightarrow Ag}$$

Anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

NO₃⁻ remains in solution while H⁺ is generated.

Thus HNO₃ effectively accumulates:$$\boxed{[\mathrm{HNO_3}]\uparrow}$$

---

M. Inert vs active electrodes

62. CuSO₄ with Cu electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

Therefore:$$\boxed{\text{Cu deposition at cathode; Cu dissolution at anode}}$$

---

63. CuSO₄ with Pt electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Therefore:$$\boxed{\mathrm{Cu+O_2}}$$

---

64. Why does anode product change?

Cu is an active electrode and can itself undergo oxidation:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

Pt is essentially inert under these conditions, so water is oxidized:$$\mathrm{H_2O\rightarrow O_2}$$

Hence:$$\boxed{\text{Nature of electrode determines the anodic reaction}}$$

---

65. Which electrode can participate?

Pt and graphite are generally inert.

Cu can participate:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

Therefore the intended answer is:$$\boxed{\text{Cu / active electrode}}$$

Answer: D if “both B and sometimes other active electrodes” is the intended wording.

---

66. CuCl₂ with Cu electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{Cu\rightarrow Cu^{2+}+2e^-}$$

Therefore:$$\boxed{\mathrm{Cu\ deposited\ at\ cathode,\ Cu\ dissolves\ at\ anode}}$$

---

67. CuCl₂ with Pt electrodes

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Therefore:$$\boxed{\mathrm{Cu+Cl_2}}$$

---

N. JEE Advanced-type questions

68. NaCl + CuSO₄ solution, inert electrodes

Ions present:$$Na^+,Cu^{2+},Cl^-,SO_4^{2-}$$

At cathode:

Cu²⁺ is reduced preferentially:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

After Cu²⁺ becomes depleted, water can produce H₂.

At anode:

In sufficiently concentrated chloride solution:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

With decreasing chloride concentration, oxygen evolution becomes increasingly important.

So the important sequence is:$$\boxed{\text{Cathode: Cu → H₂}}$$$$\boxed{\text{Anode: Cl₂ → potentially O₂}}$$

---

69. Cu²⁺ and Cl⁻ solution

At cathode:$$\mathrm{Cu^{2+}\rightarrow Cu}$$

At anode, Cl⁻ can be oxidized:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Water can also be oxidized:$$\mathrm{2H_2O\rightarrow O_2+4H^++4e^-}$$

Which dominates depends on concentration and electrode conditions.$$\boxed{\mathrm{Cl_2\ or\ O_2}}$$

can therefore be possible anodic products.

---

70. Concentrated CuCl₂ + Pt

Cathode:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

Anode:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Therefore:$$\boxed{\mathrm{Cu\ at\ cathode,\ Cl_2\ at\ anode}}$$

---

71. Dilute CuCl₂ + Pt

Cathode:$$\boxed{\mathrm{Cu}}$$

At anode, Cl₂ formation competes with O₂ evolution.

For sufficiently dilute chloride:$$\boxed{\mathrm{O_2\ can\ become\ the\ major\ anodic\ product}}$$

This is a classic concentration-based JEE concept.

---

72. Aqueous Na₂SO₄ + Pt

Na⁺ and SO₄²⁻ are essentially spectator ions.

Cathode:$$\mathrm{H_2O\rightarrow H_2}$$

Anode:$$\mathrm{H_2O\rightarrow O_2}$$

Thus Na⁺ and SO₄²⁻ are not consumed in the ideal overall reaction.$$\boxed{\text{Their amounts remain essentially unchanged}}$$

However, because water is consumed, the concentration of the dissolved Na₂SO₄ can increase as electrolysis proceeds.

---

73. Why aren’t Na⁺ and OH⁻ both discharged in aqueous NaOH?

At cathode, Na⁺ is not preferentially reduced.

Instead:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

At anode:$$\mathrm{4OH^-\rightarrow O_2+2H_2O+4e^-}$$

Therefore:$$\boxed{H_2\text{ at cathode and }O_2\text{ at anode}}$$

---

74. Fe³⁺, Cu²⁺ and H⁺

Fe³⁺ can be reduced to Fe²⁺:$$\mathrm{Fe^{3+}+e^-\rightarrow Fe^{2+}}$$

This occurs much more readily than reduction of Cu²⁺ to Cu under ordinary conditions.

Thus:$$\boxed{\mathrm{Fe^{3+}\rightarrow Fe^{2+}}}$$

Important: Fe³⁺ does not mean iron metal is deposited first. The first reduction step is Fe³⁺ → Fe²⁺.

---

75. Cl⁻, OH⁻ and SO₄²⁻

At an inert anode, chloride can be oxidized preferentially under suitable/concentrated conditions:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

OH⁻/water can produce O₂.

Therefore, under typical chloride-containing suitable conditions:$$\boxed{\mathrm{Cl^-\rightarrow Cl_2}}$$

But concentration and electrode conditions matter.

---

76. Aqueous CaCl₂ + inert electrodes

At cathode, Ca²⁺ is not normally deposited from aqueous solution.

Water is reduced:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

At anode, in sufficiently concentrated chloride solution:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Therefore:$$\boxed{\mathrm{H_2+Cl_2}}$$

and Ca²⁺ combines with OH⁻ in the vicinity of the cathode:$$\mathrm{Ca^{2+}+2OH^-\rightarrow Ca(OH)_2}$$

---

77. Aqueous MgCl₂ + inert electrodes

Cathode:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Anode, especially concentrated chloride:$$\mathrm{2Cl^-\rightarrow Cl_2+2e^-}$$

Thus:$$\boxed{\mathrm{H_2+Cl_2}}$$

Mg(OH)₂ can precipitate near the cathode:$$\mathrm{Mg^{2+}+2OH^-\rightarrow Mg(OH)_2}$$

---

78. Why isn’t Mg²⁺ deposited from aqueous MgCl₂?

Mg²⁺ requires very negative reduction conditions:$$\mathrm{Mg^{2+}+2e^-\rightarrow Mg}$$

In aqueous solution, water is reduced preferentially:$$\mathrm{2H_2O+2e^-\rightarrow H_2+2OH^-}$$

Hence:$$\boxed{\mathrm{H_2\ forms\ instead\ of\ Mg}}$$

To obtain Mg metal, molten MgCl₂ is used.

---

79. Why can Cu²⁺ be deposited but Na⁺ cannot?

Cu²⁺ has a relatively high reduction potential:$$\mathrm{Cu^{2+}+2e^-\rightarrow Cu}$$

so it can compete successfully with water.

Na⁺ has a very negative reduction potential:$$\mathrm{Na^++e^-\rightarrow Na}$$

so water is reduced instead.

Thus:$$\boxed{\mathrm{Cu^{2+}\rightarrow Cu}}$$

but:$$\boxed{\mathrm{Na^+\not\rightarrow Na\ in\ ordinary\ aqueous\ electrolysis}}$$

---

80. Na⁺, H⁺, Cl⁻ and OH⁻

At cathode, reduction occurs.

Between Na⁺ and H₂O/H⁺, hydrogen formation is preferred:$$\boxed{\mathrm{H^+ / H_2O\rightarrow H_2}}$$

At anode, oxidation occurs.

Cl⁻ is preferentially oxidized under suitable chloride concentration:$$\boxed{\mathrm{Cl^-\rightarrow Cl_2}}$$

So the common JEE answer is:$$\boxed{\text{Cathode: H₂}}$$$$\boxed{\text{Anode: Cl₂}}$$

---

🔥 JEE MASTER TABLE — MEMORISE THIS

Electrolyte	Electrode	Main product
Molten NaCl	Cathode	Na
	Anode	Cl₂
Aqueous NaCl (concentrated)	Cathode	H₂
	Anode	Cl₂
Aqueous NaCl (dilute)	Cathode	H₂
	Anode	O₂ can predominate
Molten CaCl₂	Cathode	Ca
	Anode	Cl₂
Aqueous CaCl₂	Cathode	H₂
	Anode	Cl₂ under suitable conditions
Aqueous MgCl₂	Cathode	H₂
	Anode	Cl₂ under suitable conditions
Molten Al₂O₃	Cathode	Al
	Anode	O₂ with inert anode
CuSO₄ + Cu electrodes	Cathode	Cu
	Anode	Cu dissolves
CuSO₄ + Pt electrodes	Cathode	Cu
	Anode	O₂
AgNO₃ + Ag electrodes	Cathode	Ag
	Anode	Ag dissolves
AgNO₃ + Pt electrodes	Cathode	Ag
	Anode	O₂
KI(aq) + inert	Cathode	H₂
	Anode	I₂
KBr(aq) + inert	Cathode	H₂
	Anode	Br₂
Na₂SO₄(aq) + inert	Cathode	H₂
	Anode	O₂
H₂SO₄(aq) + inert	Cathode	H₂
	Anode	O₂
KOH(aq) + inert	Cathode	H₂
	Anode	O₂
CuCl₂ + Pt	Cathode	Cu
	Anode	Cl₂
CuCl₂ + Cu	Cathode	Cu
	Anode	Cu dissolves

⭐ 5 rules that solve most JEE questions

Rule 1 — Molten vs aqueous

Molten salt:$$\boxed{\text{metal ion can be discharged}}$$

Aqueous solution:$$\boxed{\text{water competes with metal ion}}$$

---

Rule 2 — Active vs inert anode

Active metal anode:$$\boxed{\mathrm{M\rightarrow M^{n+}+ne^-}}$$

Inert anode: water/anion is oxidized.

---

Rule 3 — Cathode

Think:$$\boxed{\text{easily reducible cation first}}$$

Typical:$$\boxed{\mathrm{Ag^+>Cu^{2+}>H^+>Zn^{2+}>Na^+}}$$

---

Rule 4 — Anode

For halides:$$\boxed{\mathrm{I^->Br^->Cl^->F^-}}$$

So:$$\boxed{I_2>Br_2>Cl_2}$$

in ease of formation by oxidation.

---

Rule 5 — Faraday

$$\boxed{m=\frac{MIt}{nF}}$$

and for the same quantity of electricity:$$\boxed{m\propto\frac{M}{n}}$$

This single relation handles a huge number of JEE electrolysis numericals.

silicon dioxide (sio₂) — jee/neet short notes

1. silica / silicon dioxide

• SiO₂ = silica
• About 95% of earth’s crust is made up of silica and silicates.
• Important crystalline forms:
  • quartz
  • cristobalite
  • tridymite
• These forms are interconvertible at suitable temperatures.

2. structure of sio₂

• SiO₂ is a covalent, 3-dimensional network solid.
• Each Si atom is tetrahedrally bonded to 4 O atoms.
• Each oxygen atom is bonded to 2 silicon atoms.
• Each corner of a SiO₄ tetrahedron is shared with another tetrahedron.
• Forms 8-membered rings containing alternate Si and O atoms.
• Can be considered a giant molecule.

3. chemical properties

• SiO₂ is almost non-reactive because of its very high Si–O bond enthalpy.
• Resists:
  • halogens
  • H₂
  • most acids
  • metals, even at high temperature
• Exceptions: HF and NaOH

With NaOH:$$SiO_2+2NaOH\rightarrow Na_2SiO_3+H_2O$$

With HF:$$SiO_2+4HF\rightarrow SiF_4+2H_2O$$

4. uses of silica

• Quartz: piezoelectric material → accurate clocks, radio/TV broadcasting and mobile communications.
• Silica gel: drying agent; support for chromatographic materials and catalysts.
• Kieselguhr: amorphous silica; used in filtration.

---

silicones

• Silicones are organosilicon polymers.
• Repeating unit:

$$\boxed{(R_2SiO)_n}$$

• Starting materials: alkyl/aryl substituted silicon chlorides

$$R_nSiCl_{4-n}$$

• Methyl chloride reacts with silicon in presence of Cu catalyst at 573 K to produce methyl-substituted chlorosilanes, which are used to make silicones.

properties of silicones

• Water-repellent (hydrophobic) due to non-polar alkyl groups.
• High thermal stability.
• High dielectric strength.
• Resistant to oxidation and chemicals.
• Generally chemically stable.

uses

• Sealants
• Greases
• Electrical insulators
• Waterproofing fabrics
• Surgical and cosmetic applications because of biocompatibility

---

silicates

• Basic structural unit:

$$\boxed{SiO_4^{4-}}$$

• Si is bonded tetrahedrally to four O atoms.
• SiO₄ units can share 1, 2, 3 or 4 oxygen atoms.
• Depending on sharing, they form:
  • chains
  • rings
  • sheets
  • 3-D networks
• Negative charge is balanced by positively charged metal ions (cations).

examples

• Feldspar
• Zeolites
• Mica
• Asbestos

Important: Glass and cement are important man-made silicates.

---

zeolites

• When some Si atoms are replaced by Al atoms in the 3-D SiO₂ network, an aluminosilicate is formed.
• Replacement produces an overall negative charge.
• Charge is balanced by cations such as:
  • Na⁺
  • K⁺
  • Ca²⁺

important uses

1. Catalysts
   • Used in petrochemical industries for:
     • cracking of hydrocarbons
     • isomerisation
   • ZSM-5 converts alcohols directly into gasoline.
2. Ion exchangers
   • Hydrated zeolites are used to soften hard water by ion exchange.

JEE/NEET CHEMISTRY — 50 MARKS

Time: 60 Minutes
Maximum Marks: 50

SECTION A — MCQs

20 × 1 = 20 marks

1. Which of the following is the correct order of acidic strength in aqueous medium?
A. H₂O > CH₃OH > NH₃
B. CH₃OH > H₂O > NH₃
C. NH₃ > H₂O > CH₃OH
D. H₂O > NH₃ > CH₃OH

2. The noble gas configuration of phosphorus (Z = 15) is:
A. [Ne] 3s² 3p³
B. [He] 2s² 2p⁶ 3s² 3p³
C. [Ar] 3s² 3p³
D. [Ne] 3s² 3p⁵

3. An element having electronic configuration [Ne] 3s² 3p⁵ belongs to:
A. Group 15, Period 3
B. Group 17, Period 3
C. Group 17, Period 2
D. Group 7, Period 3

4. Which group shows the strongest +R effect?
A. –NO₂
B. –COOH
C. –NH₂
D. –CN

5. The reagents used in the Gattermann–Koch reaction are:
A. CO + HCl/AlCl₃, CuCl
B. CO₂ + HCl/AlCl₃
C. CO + Cl₂/FeCl₃
D. HCN + HCl/AlCl₃

6. Which ligand is strongest according to the spectrochemical series?
A. Cl⁻
B. F⁻
C. NH₃
D. CN⁻

7. Reimer–Tiemann reaction of phenol with CHCl₃/NaOH gives mainly:
A. Benzaldehyde
B. Salicylaldehyde
C. Benzoic acid
D. Toluene

8. The major product of Kolbe’s reaction of sodium phenoxide followed by acidification is:
A. Benzoic acid
B. Salicylic acid
C. Phenol
D. Benzaldehyde

9. A secondary amine reacts with nitrous acid to form:
A. Nitroalkane
B. N-nitrosoamine
C. Alcohol
D. Alkene

10. The equivalent weight of H₂SO₄ in an acid-base reaction is:
A. 98
B. 49
C. 24.5
D. 196

11. For a hydrogen-like atom, the energy of an electron in orbit n is proportional to:
A. n²
B. 1/n
C. –1/n²
D. –n²

12. The correct order of dipole moment is:
A. HF > HCl > HBr > HI
B. HI > HBr > HCl > HF
C. HCl > HF > HBr > HI
D. HF > HI > HBr > HCl

13. A reaction in which two species undergo oxidation-state changes to produce a species having an intermediate oxidation state is called:
A. Disproportionation
B. Comproportionation
C. Neutralisation
D. Hydrolysis

14. Which pair represents structural isomers?
A. cis-2-butene and trans-2-butene
B. n-butane and isobutane
C. (R)-2-butanol and (S)-2-butanol
D. cis-2-butene and 2-butene

15. Which alkane conformation is the most stable?
A. Eclipsed
B. Fully eclipsed
C. Staggered
D. Partially eclipsed

16. Ozonolysis of propene followed by reductive work-up gives:
A. Methanal + ethanal
B. Methanal + methanol
C. Ethanal + ethanol
D. Acetone + methanal

17. Which statement is correct for an SN2 reaction?
A. It proceeds through carbocation formation
B. Rate depends only on substrate concentration
C. It involves backside attack
D. It is favoured by tertiary alkyl halides

18. Acid chlorides react with water to produce:
A. Alcohol
B. Aldehyde
C. Carboxylic acid
D. Ether

19. Hydration of a terminal alkyne in the presence of HgSO₄/H₂SO₄ ultimately gives:
A. Primary alcohol
B. Aldehyde
C. Ketone
D. Ether

20. The reaction of an alkene with cold, dilute alkaline KMnO₄ is commonly used as:
A. Baeyer test
B. Tollens’ test
C. Fehling’s test
D. Lucas test

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SECTION B — STATEMENT / CONCEPT

10 × 2 = 20 marks

21. Consider the following statements about SN1 and SN2 reactions:

1. SN1 reactions involve carbocation formation.
2. SN2 reactions proceed through a single-step mechanism.
3. SN1 is generally favoured by tertiary substrates.
4. SN2 reactions show retention of configuration.

Which statements are correct?
A. 1, 2 and 3 only
B. 1 and 4 only
C. 2 and 3 only
D. All four

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22. For a zero-order reaction, the time required for complete consumption of reactant is finite, whereas for a first-order reaction, complete consumption theoretically requires infinite time.

If the half-life of a first-order reaction is 20 min, the time required for 75% completion is:

A. 20 min
B. 30 min
C. 40 min
D. 60 min

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23. HBr adds to propene in the presence of peroxide.

Answer the following:

(i) State the major product.
(ii) Name the rule responsible for this orientation.
(iii) Why does this reaction proceed through a radical mechanism?

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24. Match the following:

Reaction	Product/feature
A. Clemmensen reduction	1. Zn(Hg)/HCl
B. Wolff–Kishner reduction	2. NH₂NH₂/KOH
C. Cannizzaro reaction	3. Aldehydes without α-H
D. Aldol reaction	4. Aldehydes/ketones having α-H

Choose the correct matching.

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25. Give the products obtained in the following tests:

(i) Alkene + Br₂/CCl₄
(ii) Terminal alkyne + ammoniacal AgNO₃
(iii) Alkene + cold alkaline KMnO₄

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26. A Grignard reagent reacts with a nitrile followed by hydrolysis.

(i) Write the general reaction.
(ii) What class of compound is obtained?
(iii) What happens to the nitrile carbon?

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27. Complete the following reactions:

(i) CH₃CH₂Br + alc. KOH → ?
(ii) CH₃CH₂Br + aq. KOH → ?
(iii) CH₃CH₂OH —conc. H₂SO₄, 443 K→ ?

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28. Complete the following reactions:

(i) 2CH₃CHO + dilute NaOH → ?
(ii) 2C₆H₅CHO + concentrated NaOH → ?

Name both reactions.

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29. An ideal gas expands from V₁ to V₂.

Write the expression for work done when expansion is:

(i) Reversible and isothermal
(ii) Against constant external pressure

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30. Define the following:

(i) Heat capacity
(ii) Specific heat capacity
(iii) Molar heat capacity

Also write the relationship between heat supplied and temperature change.

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SECTION C — NUMERICAL / JEE-LEVEL

5 × 2 = 10 marks

31. The enthalpy change of a reaction at constant pressure is –50 kJ mol⁻¹. If the reaction involves gaseous reactants and products, explain why ΔH and ΔU may differ. Write the relation between ΔH and ΔU for gaseous reactions.

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32. Calculate the heat required to raise the temperature of 200 g of water from 25°C to 75°C, given:

Specific heat of water = 4.2 J g⁻¹ K⁻¹.

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33. The enthalpy of neutralisation of a strong acid by a strong base is approximately:

A. +57.1 kJ mol⁻¹
B. –57.1 kJ mol⁻¹
C. –5.71 kJ mol⁻¹
D. +5.71 kJ mol⁻¹

Explain the reason for its nearly constant value.

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34. Using Hess’s law:

C + ½O₂ → CO  ΔH = –110.5 kJ mol⁻¹
CO + ½O₂ → CO₂ ΔH = –283.0 kJ mol⁻¹

Calculate ΔH for:

C + O₂ → CO₂

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35. For the reaction:

N₂(g) + 3H₂(g) → 2NH₃(g)

Predict the sign of ΔS and explain your answer.

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BONUS JEE/NEET REACTION REVISION

Before attempting the paper, revise these high-yield reaction pairs:

• Reimer–Tiemann: Phenol → salicylaldehyde
• Kolbe–Schmitt: Phenol → salicylic acid
• Gattermann–Koch: Benzene → benzaldehyde
• Clemmensen: C=O → CH₂ using Zn(Hg)/HCl
• Wolff–Kishner: C=O → CH₂ using NH₂NH₂/KOH
• Cannizzaro: Aldehyde without α-H → alcohol + carboxylate
• Aldol: Carbonyl compound with α-H → β-hydroxy carbonyl compound
• Ozonolysis: Alkene → carbonyl compounds
• Baeyer test: Alkene/alkyne decolourises alkaline KMnO₄
• Ammoniacal AgNO₃: Terminal alkyne gives silver acetylide
• Br₂/CCl₄: Unsaturation test
• Grignard + nitrile: → ketone after hydrolysis
• Nitrile hydrolysis: → carboxylic acid under complete hydrolysis

ans of all only ans

ANSWER KEY — 50 MARKS

1. B
2. A
3. B
4. C
5. A
6. D
7. B
8. B
9. B
10. B
11. C
12. A
13. B
14. B
15. C
16. A
17. C
18. C
19. C
20. A
21. A
22. C
23. 1-bromopropane; Anti-Markovnikov rule; free-radical mechanism
24. A–1, B–2, C–3, D–4
25. (i) Vicinal dibromoalkane (ii) Silver acetylide precipitate (iii) Vicinal diol
26. RMgX + R′CN → R′C(=N MgX)R →[H₃O⁺] R′COR; ketone; nitrile carbon becomes carbonyl carbon
27. (i) CH₂=CH₂ (ii) CH₃CH₂OH (iii) CH₂=CH₂
28. (i) 3-hydroxybutanal (aldol) (ii) C₆H₅CH₂OH + C₆H₅COONa (Cannizzaro)
29. (i) w = −nRT ln(V₂/V₁) (ii) w = −Pext(V₂−V₁)
30. Heat capacity = q/ΔT; Specific heat = q/(mΔT); Molar heat capacity = q/(nΔT)
31. ΔH = ΔU + Δn(g)RT
32. 42,000 J = 42 kJ
33. B
34. −393.5 kJ mol⁻¹
35. ΔS < 0

1. Pathogens

• Disease-causing organisms are called pathogens.
• They include bacteria, viruses, fungi, protozoans and helminths.
• Pathogens enter the body, multiply and interfere with normal body functions.

2. Typhoid

• Pathogen: Salmonella typhi (bacterium)
• Transmission: Contaminated food and water
• Symptoms: High fever (39–40°C), weakness, stomach pain, constipation, headache and loss of appetite.
• Test: Widal test
• Severe cases may cause intestinal perforation and death.

3. Pneumonia

• Pathogens: Streptococcus pneumoniae, Haemophilus influenzae
• Affected organ: Alveoli of lungs
• Transmission: Respiratory droplets/aerosols and contaminated utensils.
• Symptoms: Fever, chills, cough and headache.
• Fluid accumulates in alveoli, causing breathing difficulty.

4. Common Cold

• Pathogen: Rhinoviruses
• Affected: Nose and respiratory passage, not lungs.
• Symptoms: Nasal congestion, nasal discharge, sore throat, cough, headache and tiredness.
• Usually lasts 3–7 days.
• Spread through droplets and contaminated objects.

5. Malaria

• Pathogen: Plasmodium
• Species include P. vivax, P. malariae and P. falciparum.
• Most serious: P. falciparum → malignant malaria.
• Vector: Female Anopheles mosquito.
• Infective form entering humans: Sporozoites.
• Parasites multiply first in liver cells, then infect RBCs.
• RBC rupture releases haemozoin, causing recurring chills and fever.
• Plasmodium requires two hosts: humans and mosquitoes.

6. Amoebiasis

• Pathogen: Entamoeba histolytica
• Affected: Large intestine
• Symptoms: Constipation, abdominal pain/cramps and stools containing excess mucus and blood.
• Transmission: Contaminated food and water.
• Carrier: Housefly acts as a mechanical carrier.

7. Ascariasis

• Pathogen: Ascaris (roundworm)
• Symptoms: Internal bleeding, muscular pain, fever, anaemia and intestinal blockage.
• Transmission: Contaminated food, water, vegetables and fruits.

8. Filariasis / Elephantiasis

• Pathogen: Wuchereria bancrofti and W. malayi
• Vector: Female mosquito
• Causes chronic inflammation of lymphatic vessels, especially of lower limbs.
• May also affect genital organs and cause deformities.

9. Ringworm

• Fungi: Microsporum, Trichophyton, Epidermophyton
• Affected: Skin, nails and scalp.
• Symptoms: Dry, scaly lesions and intense itching.
• Favourable conditions: Heat and moisture.
• Spread through infected soil, towels, clothes and combs.

10. Prevention and Control

Personal hygiene:

• Keep body clean.
• Drink clean water.
• Eat clean and properly prepared food, vegetables and fruits.

Public hygiene:

• Proper disposal of waste and excreta.
• Regular cleaning and disinfection of water reservoirs and public places.

Vector control:

• Avoid stagnant water.
• Use mosquito nets.
• Clean coolers regularly.
• Use wire mesh on doors/windows.
• Use insecticides where appropriate.
• Gambusia fish feed on mosquito larvae.

11. Vaccination and Antibiotics

• Vaccination and immunisation have helped control many infectious diseases.
• Smallpox has been completely eradicated.
• Vaccines have greatly controlled diseases such as polio, diphtheria, pneumonia and tetanus.
• Antibiotics and other drugs help in treating many infectious diseases.