NCERT solutions for class 12 chemistry chapter 3 – Electrochemistry is now available for free download. The solutions are given in a PDF format which will make it easier for students to access and refer to it. NCERT Solutions chemistry chapter 3 includes various questions on the the topic of electrochemistry. These solutions have been designed to help students understand and get used to all the concepts of the chapter. Besides, students can practice questions from these materials as the questions are purely based on NCERT textbooks that is prescribed for class 12 in CBSE schools. The solutions are prepared by the best subject experts. In essence, this NCERT solutions can be useful for those students preparing for class 12 board exams and also for JEE advance and other medical entrance exams.

The NCERT solutions for chapter 3 – electrochemistry has mainly been designed to help the students help them prepare well and score good marks in CBSE class 12 chemistry paper. Further, the solutions consist of well thought of or structured questions along with detailed explanations to help students learn and remember concepts easily. You have come to the right place as you will be getting access to detailed, accurate and free solutions for class 12 NCERT chemistry. The solutions can be viewed either online on the website or you can download the PDF for later viewing without the need for any internet connection.

*Q 3.1:*

*In the order of their reactivity, i.e how they displace each other from their salt solutions, allign the metals in decreasing order. Cu, Fe, Al, Zn and Mg.*

*Answer:*

According to their reactivity, the given metals replace the others from their salt solutions in the said order: Mg, Al, Zn, Fe, Cu .

Mg : Al : Zn : Fe : Cu

* *

*Q 3.2:*

*Standard electrode potentials given as,*

*Mg ^{2+}/Mg = −2.37 V, Hg^{2+}/Hg = 0.79V, Cr^{3+}/Cr = − 0.74V, Ag^{+}/Ag = 0.80V, K^{+}/K = −2.93V*

*In the order of increasing of reducing power arrange the given metals accordingly.*

*Ans:*

The reducing power increases with the lowering of reduction potential. In order of given standard electrode potential (increasing order) : K^{+}/K < Mg^{2+}/Mg < Cr^{3+}/Cr < Hg^{2+}/Hg < Ag^{+}/Ag

Thus, in the order of reducing power, we can arrange the given metals as : Ag< Hg < Cr < Mg < K

* *

*Q 3.3 : *

*Represent the galvanic cell in which the following reaction takes place.*

*Zn(s) + 2Ag ^{+}(aq) → Zn^{2+}(aq) + 2Ag(s) *

*Also find :*

*(i) The negatively charged electrode ?*

*(ii) Current carriers in the cell.*

*(iii) At each electrode, the individual reaction.*

*Ans :*

The galvanic cell in which the given reaction takes place is depicted as:

\(Zn_{ ( s ) } | Zn^{ 2+ }_{( aq )}||Ag^{ + }_{( aq )}|Ag_{( s )}\)

(i) The negatively charged electrode is the Zn electrode (anode)

(ii) The current carriers in the cell are ions. Current flows to zinc from silver in the external circuit.

(iii) Reaction at the anode is given by :

\(Zn_{ ( s ) }\rightarrow Zn^{ 2+ }_{( aq )} + 2 e^-\)

Reaction at the anode is given by :

\(Ag^{+}_{ ( aq ) } + e^- \rightarrow Ag_{( s )}\)

* *

*Q 3.4:*

*With the following reactions given, find the standard cell potentials of galvanic cells with given reactions.*

*(i) 2Cr _{(s)} + 3Cd^{2+}_{(aq)} → 2Cr^{3+}_{(aq)} + 3Cd*

*(ii) Fe ^{2+}_{(aq)} + Ag^{+}_{(aq)} → Fe^{3+}_{(aq)} + Ag_{(s)}*

*Calculate the ∆ _{r}G^{θ} and equilibrium constant of the reactions.*

*Ans :*

(i) \(E^{\Theta}_{Cr^{3+}/Cr}\) = 0.74 V

\(E^{\Theta}_{Cd^{2+}/Cd}\) = -0.40 V

The galvanic cell of the given reaction is depicted as :

\(Cr_{ ( s ) }|Cr^{ 3+ }_{ ( aq ) }||Cd^{ 2+ }_{ aq }|Cd_{ ( s ) }\)

Now, the standard cell potential is

\(E^{\Theta }_{cell} = E^{\Theta }_{g}-E^{\Theta }_{L}\)

= – 0.40 – ( -0.74 )

= + 0.34 V

In the given equation, n = 6

F = 96487 C mol^{−1}

\(E^{\Theta }_{cell}\) = + 0.34 V

Then, \(\Delta_rG^{\Theta}\) = −6 × 96487 C mol^{−1} × 0.34 V

= −196833.48 CV mol−1

= −196833.48 J mol−1

= −196.83 kJ mol−1

Again,

\(\Delta_rG^{\Theta} = – R T ln K\)

\(\Delta_rG^{\Theta} = – 2.303 R T ln K\)

\(log {k} = \frac{\Delta_rG}{ 2.303 R T }\)

\(= \frac{-196.83\times10^{3}}{ 2.303 \times8.314\times 298 }\)

= 34.496

K = antilog (34.496) = 3.13 × 10^{34}

The galvanic cell of the given reaction is depicted as:

\(Fe^{ 2+ }_{( aq )}|Fe^{ 3+ }_{( aq )}|| Ag^+_{( aq )}|Ag_{( s )}\)

Now, the standard cell potential is

\(E^{\Theta }_{cell} = E^{\Theta }_{g}-E^{\Theta }_{L}\)

Here, n = 1.

Then, \(\Delta_t G^0 = -nFE^0_{cell}\)

= −1 × 96487 C mol^{−1} × 0.03 V

= −2894.61 J mol^{−1}

= −2.89 kJ mol^{−1}

Again, \(\Delta_t G^0 = -2.303 RT \; ln K\)

\(ln K = \frac{\Delta_t G}{ 2.303 RT }\)

\(= \frac{-2894.61 }{ 2.303 \times 8.314 \times 298 }\)

= 0.5073

K = antilog (0.5073)

= 3.2 (approximately)

* *

*Q 3.5:*

*Write the Nernst equation and emf of the following cells at 298 K:*

*(i) Mg _{(s)} | Mg^{2+}(0.001M) || Cu^{2+}(0.0001 M) | Cu_{(s)}*

*(ii) Fe _{(s)} | Fe^{2+}(0.001M) || H^{+}(1M)|H_{2(g)}(1bar) | Pt_{(s)} *

*(iii) Sn _{(s)} | Sn^{2+}(0.050 M) || H^{+}(0.020 M) | H_{2(g) }(1 bar) | Pt_{(s)}*

*(iv) Pt _{(s)} | Br_{2(l)} | Br^{−}(0.010 M) || H^{+}(0.030 M) | H_{2(g)} (1 bar) | Pt_{(s)}.*

*Answer*

(i) For the given reaction, the Nernst equation can be given as:

\(E_{cell} = E^0_{cell} – \frac{0.591}{n}log\frac{[Mg^{2+}]}{[Cu^{2+}]}\)

\(= 0.34 – (-2.36) – \frac{0.0591}{2} log \frac{0.001}{0.0001}\)

\(2.7 -\frac{0.0591}{2}log10\)

= 2.7 − 0.02955

= 2.67 V (approximately)

(ii) For the given reaction, the Nernst equation can be given as:

\(E_{cell} = E^0_{cell} – \frac{0.591}{n}log\frac{[Fe^{2+}]}{[H ^{+}]^2}\)

= 0 – ( – 0.14) – \(\frac{0.0591}{n}log\frac{0.050}{(0.020)^{2}}\)

= 0.52865 V

= 0.53 V (approximately)

(iii) For the given reaction, the Nernst equation can be given as:

\(E_{cell} = E^0_{cell} – \frac{0.591}{n}log\frac{[Sn^{2+}]}{[H ^{+}]^2}\)

= 0 – ( – 0.14) – \(\frac{0.591}{2}log\frac{0.050}{(0.020)^2}\)

= 0.14 − 0.0295 × log125

= 0.14 − 0.062

= 0.078 V

= 0.08 V (approximately)

(iv) For the given reaction, the Nernst equation can be given as:

\(E_{cell} = E^0_{cell} – \frac{0.591}{n}log\frac{1}{[Br^{-}]^2[H ^{+}]^2}\)

= 0 – 1.09 – \(\frac{0.591}{2}log\frac{1}{(0.010)^2(0.030)^2}\)

= -1.09 – 0.02955 x \(log\frac{1}{0.00000009}\)

= -1.09 – 0.02955 x \(log\frac{1}{9\times 10^{-8}}\)

= -1.09 – 0.02955 x \(log{ (1.11 \times 10^{7} )}\)

= -1.09 – 0.02955 x (0.0453 + 7)

= -1.09 – 0.208

= -1.298 V

* *

*Q 3.6:*

*The following reaction takes place in the button cells widely used in watches and other devices:*

* For the given reaction calculate \(\Delta_r G^\Theta\) and \(E^0\) : *

*Ans: *

\(E^0\) = 1.104 V

We know that,

\(\Delta_r G^\Theta = -nFE^\Theta\)

= −2 × 96487 × 1.04

= −213043.296 J

= −213.04 kJ

*Q 3.7:*

*For the solution of an electrolyte describe its conductivity and molar conductivity. Also put some light on how they vary with concentration.*

*Answer*

Conductivity of a solution is defined as the conductance of a solution of 1 cm in length and area of cross – section 1 sq. cm. Specific conductance is the inverse of resistivity and it is represented by the symbol κ. If ρ is resistivity, then we can write:

\(k = \frac{1}{\rho}\)

At any given concentration, the conductivity of a solution is defined as the unit volume of solution kept between two platinum electrodes with the unit area of cross- section at a distance of unit length.

\(G = k \frac{a}{l} = k \times 1 = k\) [Since a = 1 , l = 1]

When concentration decreases there will a decrease in Conductivity. It is applicable for both weak and strong electrolyte. This is because the number of ions per unit volume that carry the current in a solution decreases with a decrease in concentration.

**Molar conductivity – **

Molar conductivity of a solution at a given concentration is the conductance of volume V of a solution containing 1 mole of the electrolyte kept between two electrodes with the area of cross-section A and distance of unit length.

\(\Lambda_m = k \frac{A}{l}\)

Now, l = 1 and A = V (volume containing 1 mole of the electrolyte).

\(\Lambda_m = k V\)

Molar conductivity increases with a decrease in concentration. This is because the total volume V of the solution containing one mole of the electrolyte increases on dilution. The variation of \(\Lambda_m\) with \(\sqrt{c}\) for strong and weak electrolytes is shown in the following plot :

**Q 3.8:**

The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 Scm^{−1}. Find its molar conductivity.

Ans :

Given, κ = 0.0248 S cm^{−1} c

= 0.20 M

Molar conductivity, \(\Lambda_m = \frac{k \times 1000}{c}\)

\(= \frac{0.0248 \times 1000}{0.2}\)

= 124 Scm^{2}mol^{-1}

*Q 3.9:*

*Considering the case of a conductivity cell having 0.001 M KCl solution at 298 K is 1500 Ω. If given, conductivity of 0.001M KCl solution at 298 K is 0.146 × 10 ^{−3} S, find the cell constant?*

*Answer*

Given,

Conductivity, k = 0.146 × 10^{−3} S cm−1

Resistance, R = 1500 Ω

Cell constant = k × R

= 0.146 × 10^{−3} × 1500

= 0.219 cm^{−1}

*Q 3.10:*

*The conductivity of NaCl at 298 K has been found at different concentrations and the results are given below:*

*Concentration/M 0.001 0.010 0.020 0.050 0.100*

*10 ^{2} × k/S m^{−1} 1.237 11.85 23.15 55.53 106.74*

*for all concentrations and draw a plot between \(\Lambda_m\) and c ^{1⁄2}. Find the value*

*Molar conductivity of*

*Calculate \(\Lambda_m\) of \(\Lambda^0_m\)*

*Ans:*

Given,

κ = 1.237 × 10^{−2} S m−1, c = 0.001 M

Then, κ = 1.237 × 10^{−4} S cm^{−1}, c^{1⁄2} = 0.0316 M^{1/2}

\(\Lambda_m =\frac{k}{c}\)

\(=\frac{1.237 \times 10^{ -4 } S\;cm^{-1} }{0.001 \; mol\; \; L^{ -1 }}\times \frac{1000\;cm^{-1}}{L}\)

= 123.7 S cm^{2} mol^{−1}

Given,

κ = 11.85 × 10^{−2} S m^{−1}, c = 0.010M

Then, κ = 11.85 × 10^{−4} S cm^{−1}, c^{1⁄2} = 0.1 M^{1/2}

\(\Lambda_m =\frac{k}{c}\)

\(=\frac{11.85 \times 10^{ -4 } S\;cm^{-1} }{0.010 \; mol\; \; L^{ -1 }}\times \frac{1000\;cm^{-1}}{L}\)

= 118.5 S cm^{2} mol^{−1}

Given,

κ = 23.15 × 10^{−2} S m^{−1}, c = 0.020 M

Then, κ = 23.15 × 10^{−4} S cm^{−1}, c^{1/2} = 0.1414 M^{1/2}

\(\Lambda_m =\frac{k}{c}\)

\(=\frac{23.15 \times 10^{ -4 } S\;cm^{-1} }{0.020 \; mol\; \; L^{ -1 }}\times \frac{1000\;cm^{-1}}{L}\)

= 115.8 S cm^{2} mol^{−1 }

Given,

κ = 55.53 × 10^{−2} S m^{−1}, c = 0.050 M

Then, κ = 55.53 × 10^{−4} S cm^{−1}, c^{1/2} = 0.2236 M^{1/2}

\(\Lambda_m =\frac{k}{c}\)

\(=\frac{106.74 \times 10^{ -4 } S\;cm^{-1} }{0.050 \; mol\; \; L^{ -1 }}\times \frac{1000\;cm^{-1}}{L}\)

= 111.1 1 S cm^{2} mol^{−1}

Given,

κ = 106.74 × 10^{−2} S m^{−1}, c = 0.100 M

Then, κ = 106.74 × 10^{−4} S cm^{−1}, c^{1/2} = 0.3162 M^{1/2}

\(\Lambda_m =\frac{k}{c}\)

\(=\frac{106.74 \times 10^{ -4 } S\;cm^{-1} }{0.100 \; mol\; \; L^{ -1 }}\times \frac{1000\;cm^{-1}}{L}\)

= 106.74 S cm^{2} mol^{−1}

Now, we have the following data :

Since the line interrupts \(\Lambda_m\) at 124.0 S cm^{2} mol^{−1}, \(\Lambda^0_m\) = 124.0 S cm^{2} mol^{−1}

*Q 3.11:*

*Find the molar conductivity of acetic acid if its conductivity is given to be 0.00241 M . Also, if the value of \(\Lambda^0_m\) is given to be390.5 S cm ^{2} mol^{−1}, calculate its dissociation constant?*

*Ans:*

Given, κ = 7.896 × 10^{−5} S m^{−1} c

= 0.00241 mol L^{−1}

Then, molar conductivity, \(\Lambda_m = \frac{k}{c}\)

= \(\frac{7.896 \times 10^{-5} S cm^{-1}}{0.00241 \; mol \; L^{-1}}\times \frac{1000 cm^3}{L}\)

= 32.76S cm^{2} mol^{−1}

\(\Lambda^0_m =\) 390.5 S cm^{2} mol^{−1}

Again,

\(\alpha =\frac{\Lambda_m }{\Lambda^0_m }\)

= \(= \frac{32.76 \; S\; cm^2 \; mol^{-1} }{390.5 \; S\; cm^2 \; mol^{-1} }\)

Now,

= 0.084

Dissociation constant, \(K_a = \frac{c\alpha^2}{(1-\alpha)}\)

= \(\frac{ ( 0.00241 \; mol \; L^{-1} )( 0.084 )^2}{ ( 1 – 0.084 ) }\)

= 1.86 × 10^{−5} mol L^{−1}

*Q 3.12:*

*How much charge is required for the following reductions of 1 mol of :*

*(i) Al ^{3+} to Al.*

*(ii) Cu ^{2+} to Cu.*

*(iii)\(MnO^-_4\) to Mn ^{2+}.*

*Ans : *

(i) \(Al^{3+} + 3e^- \rightarrow Al\)

Required charge = 3 F

= 3 × 96487 C

= 289461 C

(ii) \(Cu^{2+} + 2e^- \rightarrow Cu\)

Required charge = 2 F

= 2 × 96487 C

= 192974 C

(iii) \(MnO^-_4 \rightarrow Mn^{2+}\)

i.e \(Mn^{7+} + 5e^-\rightarrow Mn^{2+}\)

Required charge = 5 F

= 5 × 96487 C

= 482435 C

*Q 3.13:*

*In the terms of Faraday, how much electricity is required to produce : *

*(i) From molten CaCl _{2}, 20.0 g of Ca.*

*(ii) From molten Al _{2}O_{3,} 40.0 g of Al.*

*Ans:*

(i) From given data,

\(Ca^{2+} + 2e^- \rightarrow Ca\)

Electricity required to produce 40 g of calcium = 2 F

Therefore, electricity required to produce 20 g of calcium = (2 x 20 )/ 40 F

= 1 F

(ii) From given data,

\(Al^{3+} + 3e^- \rightarrow Al\)

Electricity required to produce 27 g of Al = 3 F

Therefore, electricity required to produce 40 g of Al = ( 3 x 40 )/27 F

= 4.44 F

* *

*Q 3.14:*

*Calculate the amount of electricity required for the oxidation of 1 mol of the following in coulombs :*

* (i) H _{2}O to O_{2}.*

*(ii)FeO to Fe _{2}O_{3}.*

*Ans :*

(i) From given data,

\(H_2O\rightarrow H_2 + \frac{1}{2}O_2\)

We can say that :

\(O^{2-}\rightarrow \frac{1}{2}O_2 + 2e^-\)

Electricity required for the oxidation of 1 mol of H_{2}O to O_{2} = 2 F

= 2 × 96487 C

= 192974 C

(ii) From given data,

\(Fe^{2+}\rightarrow Fe^{3+} + e^-\)

Electricity required for the oxidation of 1 mol of FeO to Fe_{2}O_{3} = 1 F

= 96487 C

*Q 3.15:*

*For 20 minutes, a current of 5 A is applied to between platinum electrodes to electrolyze a solution of Ni(NO _{3})_{2}. Find the amount of Ni deposited at the cathode?*

*Ans :*

Given,

Current = 5A

Time = 20 × 60 = 1200 s

Charge = current × time

= 5 × 1200

= 6000 C

According to the reaction,

\(Ni^{2+} + 2e^-\rightarrow Ni_{ (s) } + e^-\)

Nickel deposited by 2 × 96487 C = 58.71 g

Therefore, nickel deposited by 6000 C = \(\frac{58.71 \times 6000}{2 \times 96487}g\)

= 1.825 g

Hence, 1.825 g of nickel will be deposited at the cathode.

*Q 3.16:*

*Solutions of 3 electrolytic cells are ZnSO _{4}, AgNO_{3} and CuSO_{4},cells are connected in series. Of the cells, A,B,C respectively, after passing a steady current of 1.5 amperes , 1.45 g of silver was found deposited at the cathode of cell B. How much time did the current flow? What amount of zinc and copper were deposited?*

*Ans :*

According to the reaction:

\(Ag^+_{(aq)} +e^- \rightarrow Ag_{(s)}\)

i.e., 108 g of Ag is deposited by 96487 C.

Therefore, 1.45 g of Ag is deposited by = \(\frac{96487\times 1.45}{107}C\)

= 1295.43 C

Given,

Current = 1.5 A

Time = 1295.43/ 1.5 s

= 863.6 s

= 864 s

= 14.40 min

Again,

\(Cu^{2+}_{(aq)} +2e^-\rightarrow Cu_{(s)}\)

i.e., 2 × 96487 C of charge deposit = 63.5 g of Cu

Therefore, 1295.43 C of charge will deposit \(\frac{63.5 \times 1295.43}{2 \times 96487}\)

= 0.426 g of Cu

\(Zn^{2+}_{(aq)} +2e^-\rightarrow Zn_{(s)}\)

i.e., 2 × 96487 C of charge deposit = 65.4 g of Zn

Therefore, 1295.43 C of charge will deposit \(\frac{65.4 \times 1295.43}{2 \times 96487}\)

= 0.439 g of Zn

*Q 3.17:*

*Using the standard electrode potentials given in Table 3.1, predict if the reaction between*

*the following is feasible:*

*(i)*

* Fe ^{3+}(aq) and I^{−}(aq)*

*(ii) Ag ^{+} (aq) and Cu(s)*

*(iii) Fe ^{3+} (aq) and Br^{−} (aq)*

*(iv) Ag(s) and Fe ^{3+} (aq)*

*(v) Br _{2} (aq) and Fe^{2+} (aq).*

*Ans :*

(i)

(ii)

E^{0 } is positive , hence reaction is feasible.

(iii)

E^{0 } is negative , hence reaction is not feasible.

(iv)

E^{0 } is negative , hence reaction is not feasible.

(v)

E^{0 } is positive , hence reaction is feasible.

*Q 3.18:*

*Predict the products of electrolysis in each of the following :*

*(i) An aqueous solution of AgNO _{3} with silver electrodes.*

*(ii) An aqueous solution of AgNO _{3}with platinum electrodes.*

*(iii) A dilute solution of H _{2}SO_{4}with platinum electrodes.*

*(iv) An aqueous solution of CuCl _{2} with platinum electrodes.*

*Ans:*

(i) At cathode:

The following reduction reactions compete to take place at the cathode.

\(Ag^+_{(aq)}+e^- \rightarrow Ag_{(s)}\) ; E^{0 }= 0.80 V

\(H^+_{(aq)}+e^- \rightarrow \frac{1}{2}H_{2(g)}\) ;E^{0 }= 0.00 V

The reaction with a higher value of E^{0} takes place at the cathode. Therefore, deposition of silver will take place at the cathode.

At anode:

The Ag anode is attacked by \(NO^+_3\) ions. Therefore, the silver electrode at the anode dissolves in the solution to form Ag^{+}.

(ii) At cathode:

The following reduction reactions compete to take place at the cathode.

\(Ag^+_{(aq)}+e^- \rightarrow Ag_{(s)}\) ; E^{0 }= 0.80 V

\(H^+_{(aq)}+e^- \rightarrow \frac{1}{2}H_{2(g)}\) ;E^{0 }= 0.00 V

The reaction with a higher value of E^{0} takes place at the cathode. Therefore, deposition of silver will take place at the cathode.

At anode:

Since Pt electrodes are inert, the anode is not attacked by \(NO^+_3\) ions. Therefore, OH^{−} or \(NO^+_3\) ions can be oxidized at the anode. But OH^{−} ions having a lower discharge potential and get preference and decompose to liberate O_{2}.

\(OH^-\rightarrow OH + E^-\)

\(4OH^-\rightarrow 2H_2O + O_2\)

(iii) At the cathode, the following reduction reaction occurs to produce H_{2} gas.

\(H^+_{(aq)}+e^-\rightarrow \frac{1}{2}H_{2(g)}\)

At the anode, the following processes are possible.

\(2H_2O_{(l)}\rightarrow O_{2(g)} + 4H^+_{(aq)}+4e^-\) ; E^{0} = +1.23 V —–(i)

\(2SO^{2-}_{4(aq)}\rightarrow S_2O^{2-}_{6(aq)} + 2e^-\) ; E^{0} = +1.96 V —–(ii)

For dilute sulphuric acid, reaction (i) is preferred to produce O_{2} gas. But for concentrated sulphuric acid, reaction (ii) occurs.

(iv) At cathode:

The following reduction reactions compete to take place at the cathode.

\(Cu^{2+}_{(aq)}+2e^- \rightarrow Cu_{(s)}\) ; E^{0 }= 0.34 V

\(H^+_{(aq)}+e^- \rightarrow \frac{1}{2}H_{2(g)}\) ;E^{0 }= 0.00 V

The reaction with a higher value of takes place at the cathode. Therefore, deposition of copper will take place at the cathode.

At anode:

The following oxidation reactions are possible at the anode.

\(Cl^{-}_{(aq)} \rightarrow \frac{1}{2} Cl_{2(g)}+e^-\); E^{0 }= 1.36 V

\(2H_20_{(l)} \rightarrow O_{2(g)} + 4H^+_{(aq)} +e^-\); E^{0 }= +1.23 V

At the anode, the reaction with a lower value of E^{0} is preferred. But due to the over potential of oxygen, Cl^{−} gets oxidized at the anode to produce Cl_{2} gas.

### About BYJU’s NCERT Solutions

We at BYJU’s are providing a comprehensive set of NCERT solutions for students that has further been designed to offer several benefits to the users. Well, BYJU’s solutions offer the luxury to learn NCERT class 12 chemistry syllabus including chapter 3 from anywhere and within the comfort zone of the students. Students can learn, practice, revise the different chemistry chapter 9 topics right from their homes or from any place, be it at a coffee shop, at a restaurant, shopping mall, while travelling etc. The solutions are easily accessible and students can view the solutions right on the website or they can download and use BYJU’s- The Learning app for more enhanced learning experience

With these NCERT solutions for class 12, students can easily customize how they learn. So whenever there are certain difficult topics, they can usually go at a slower pace in understanding the topics and solving the chemistry papers. Apart from these NCERT solutions, BYJU’s also has the best subject experts who can guide students to learn the subject and its concepts in a more simple and easy manner. Further, if students come across any doubts or queries while going through the NCERT class 12 chemistry solutions, they can always approach BYJU’s responsive support team for clearing all their doubts.BYJU’s also keeps a track of all the progress that student’s make and offers feedback as well as counseling via periodic assessments. Besides, students can bring in all their queries regarding chemistry, as well as other subjects including physics, biology and maths.