Unit 16: Unit 16: Electrochemistry — Short Questions

Ans. When electricity is passed through aqueous solution of sodium chloride both sodium and hydrogen cations are attracted towards the cathode. As sodium lies above hydrogen in electrochemical series, hydrogen ion is discharged at cathode and chlorine gas is produced at anode. In the solution, sodium ions Na⁺ combine with OH⁻. Therefore, it will become alkaline due to formation of NaOH.

Main products

• Hydrogen gas at cathode
• Chlorine gas at anode
• Sodium hydroxide NaOH in solution.

Ans. Objectives of electroplating

There are three main objectives of electroplating.

(i) Decoration
To deposit the noble metals like gold and silver on the inexpensive metals to enhance their beauty.

(ii) Protection
To protect the metals from rusting as well as from the reaction of organic acids.

(iii) Repair
To weld the broken parts of the machinery by electrodeposition of metals. Generally copper, silver, chromium, nickel and gold are electroplated.

Ans.

Electrolytic cell A cell which converts electrical energy into chemical energy is called electrolytic cell.

Galvanic cell A cell which uses chemical energy to produce electric current is called Galvanic cell.

Ans. It is a negative electrode that is cathode through which electrons leave the external circuit and move towards positive electrode (anode).

Ans. Electrolytes conduct electricity through the movement of ions. When dissolved in water or melted electrolyte breaks into cation and anions (positive and negative ions). Cations move towards the cathode and anions move towards the anode. This movement of ions is the reason of flow of electric current.

Ans. In the electrochemical series, copper is below iron, so it can't displace iron from FeSO₄ solution. So, no reaction occurs because copper is less reactive than iron. The copper strip remains unchanged and the green color of FeSO₄ solution stays the same.

Ans. Fuel cells are environmental friendly because:

(i) Fuel cell produce zero emission of carbon dioxide because water is the only product. Fossil fuels, on the other hand, produce carbon dioxide, nitrogen oxides and sulphur oxides which are very harmful to the environment.

(ii) Less noise pollution: Fuel cells produce less noise pollution.

(iii) More sustainable: Hydrogen used in fuel cells can be obtained from renewable sources like solar or wind power

Ans. The oxidation number is an apparent charge of an atom in a compound or an ion. It may have a positive or negative ion.

Example Oxidation number O in H₂O is -2.

Ans. Oxidation number of atoms in compounds or ions helps to identify either the reaction is oxidation or reduction. Oxidation numbers are also used to write chemical formulas and names of chemical compounds.

Ans. Solution

2(Cr)+3(-2)=-2
2(Cr)+(-6)=-2
2(Cr)-6=-2
2(Cr)=-2+6
2Cr/2=4/2
Cr=+2

Ans. When writing the formulae of ionic compound containing more than one polyatomic ions the chemical formula of ions is taken in brackets and indicate the number of such ions with subscript.

Example The formula of ammonium carbonate can be written as:

The formula of ammonium carbonate is (NH₄)₂ CO₃.

Ans. The formula of magnesium nitrate will be written as:

The formula of magnesium nitrate is Mg(NO₃)₂.

Ans.

Oxidation Oxidation is defined as a process in which oxygen is either added or hydrogen is removed.

Example
Oxygen combine with other elements in the formation of their oxides.
4Fe₍g₎ + 3O₂₍g₎ → 2Fe₂O₃₍s₎

Reduction Reduction is defined as a process in which either addition of hydrogen or removal of oxygen takes place.

Example
Hydrogen is added in chlorine to give hydrogen chloride gas.
H₂₍g₎ + Cℓ₂₍g₎ → 2HCℓ₍g₎

Ans. Reduction is a process in which oxidation number of an element decreases.

Example Mn⁷⁺ + 5e⁻ → Mn²⁺

In this example oxidation number of manganese is reduced from +7 [Mn(VII)] to +2[Mn(II)]

Ans. Reducing agents The species which oxidizes other species in a chemical reaction are called reducing agent.

Example Hydrogen is removed from a compound in the following reaction.
CH₄₍g₎ + 2O₂₍g₎ → 2CO₂₍g₎ + 2H₂O₍l₎

CH₄ is being oxidized while O₂ is being reduced.
CH₄ is reducing agent in this reaction.

Ans. Redox reactions A chemical reaction which involves simultaneous oxidation and reduction is called oxidation – reduction or redox reaction.

Example H₂S + Cℓ₂ → 2HCℓ + S

Ans. Potassium permanganate oxidizes oxalic acid into carbon dioxide. During the reaction dark purple colour of potassium permanganate change to colourless.

Ans. Electrolysis When electricity is passed through fused ionic compounds or their aqueous solutions, they are decomposed into their constituent components. Such compounds are called electrolytes and the process is called electrolysis.

Ans. In the process of electrolysis, the non-spontaneous reaction takes place with the help of electrical energy. The external voltage source forces the non-spontaneous reaction by pumping electrons towards one of the electrodes.

Ans. The positive ions of the electrolyte are attracted towards the negative electrode that is cathode, and get reduced by accepting electrons from it. So reduction takes place at cathode.

Ans. The electrolysis of a concentrated solution of sodium chloride, also called brine, produces hydrogen gas at the cathode, chlorine gas at the anode and sodium hydroxide as a by-product.

Ans. At cathodes Na⁺ and H₂O more towards cathode. Water molecules or hydrogen ions pick up electrons to form hydrogen gas.

2H₂O₍ℓ₎ + 2e⁻ → H₂₍g₎ + 2OH⁻₍aq₎

At anode Cℓ⁻ ions move towards anode and gets oxidized by giving electrons.

2Cℓ⁻₍aq₎ → Cℓ₂₍g₎ + 2e⁻

Ans. The OH⁻ ions are formed around cathode due to decomposition of water. So the solution around the cathode becomes alkaline due the formation of hydroxide ions.

Ans. The situation becomes complicated in the electrolysis of aqueous metal halide due to presence of water. Water also ionizes in small amount to give hydrogen and hydroxide ions.

Ans. The elements present at the top of the electrochemical series possess more tendency to lose electron and gets oxidized.

Example Na, Li, K etc have more tendency to get oxidized as compared to other elements.

Ans. At cathode If aqueous solution of metals like nickel or copper is electrolyzed, these are two possibilities.

i. If the solution is concentrated, metal is deposited at cathode.

ii. If solution is very dilute hydrogen gas is evolved.

At in-between concentration both hydrogen and metal are obtained.

2M₍aq₎ + 2e⁻ → 2M₍s₎
2H⁺₍aq₎ + 2e⁻ → H₂₍g₎

Ans. If concentrated aqueous solution of metal halide is electrolysed, hydrogen ions are discharged at the cathode. If the metal is present higher than hydrogen in electrochemical series. If metals like silver and copper which are lower than hydrogen in electrochemical, then metal ions are discharged at cathode.

Ans. The reactions at the electrodes are:

At cathode Metal or hydride is produced.
2M⁺₍aq₎ + 2e⁻ → 2M₍g₎
2H⁺₍aq₎ + 2e⁻ → H₂₍g₎

At anode Halogen molecules are generally formed. 2X⁻₍aq₎ → X₂₍g₎ + 2e⁻

Ans. Following are the reactions at electrodes.

At cathode Molten lead metal is obtained at cathode. Pb²⁺₍l₎ + 2e⁻ → Pb₍l₎ (oxidation)

At Anode Chlorine gas is formed at anode. 2Cℓ⁻ → Cℓ₂₍g₎ + 2e⁻ (reduction)

Ans. Electrolyte must be heated upto 501°C to get molten lead chloride. Heat will render ions free to move and conduct electricity. The electrolysis must be carried out at the temperature above the melting point of lead chloride.

Ans. Inert electrodes like platinum or graphite are used for the electrolysis of dilute solution of sulphuric acid because they did not react with the electrolyte.

Ans. The electrolysis of aqueous solution of copper (II) sulphate with inert electrodes results in the decomposition of water and copper (II) sulphate.

At cathode Copper metal is deposited at cathode (inert electrode).
Cu²⁺₍aq₎ + 2e⁻ → Cu₍s₎ (Reduction)

At anode Oxygen is produced at the anode
4OH⁻ → O₂₍g₎ + 2H₂O₍l₎ + 4e⁻ (oxidation)

Ans. In the electrolysis of aqueous solution of copper sulphate using copper electrodes copper metal from copper anode start going into the solution in the form of copper ions and electrode mass starts decreasing

At cathode Cu²⁺₍aq₎ + 2e⁻ → Cu₍s₎ (Reduction)

At anode Cu₍s₎ → Cu²⁺₍aq₎ + 2e⁻ (Oxidation)

Ans. Fuel cells Fuel cells provide a method in which chemical energy is converted into electrical energy. These cells use gaseous fuels like hydrogen and oxygen to produce electricity.

Ans. At anode Hydrogen is oxidized to water. H₂₍g₎ + 2OH⁻₍aq₎ → 2H₂O₍l₎ + 2e⁻ [2]

At cathode Oxygen is reduced to hydroxide ions. O₂₍g₎ + 2H₂O₍l₎ + 4e⁻ → 4OH⁻₍aq₎

Ans. Advantages of fuel cells

i. Fuel cells produce less noise pollution.

ii. Hydrogen used in fuel cells can be obtained from renewable sources like solar or wind power. This fact renders fuel cells more sustainable

Ans. Disadvantages of fuel cells

(i) Fuel cells at present, are very expensive because the production of hydrogen by electrolysis of water incurs high cost.

(ii) Hydrogen being highly inflammable requires very special storage and transportation facilities which further increases the cost of fuel cells.

Ans. Corrosion Corrosion is a natural process due to which a metal suffers gradual decay when it interacts chemically with its environment.

Ans. Coating of the metal surface

Corrosion can be avoided by coating the metal surface with some barriers that prevent the contact of metal surface with the atmosphere. For this purpose, the metal surfaces are protected by either greasing with oil or painting.

Ans. Galvanization It is a common method to protect iron or steel from rusting. It involves the dipping of an iron sheet in a bath of molten zinc. A layer of zinc metal gets deposited on the iron sheet and prevents rusting. Tin cans are actually steel cans which have been coated with a thin layer of tin.

Ans. In sacrificial protection a more reactive metal like zinc or magnesium is used which will corrode in place of the less reactive metal like iron. Magnesium is often employed to protect steel in buried fuel tanks and pipe lines and ship hulls submerged in sea water.

Ans. There are three main objectives of electroplating.

(i) Decoration
To deposit the noble metals like gold and silver on the inexpensive metals to enhance their beauty.

(ii) Protection
To protect the metals from rusting as well as from the reaction of organic acids.

(iii) Repair
To weld the broken parts of the machinery by electrodeposition of metals. Generally copper, silver, chromium, nickel and gold are electroplated.

Ans. The electrolyte chosen should have the following properties:

(i) very soluble in water

(ii) good conductor

(iii) not easily oxidized or reduced or hydrolyzed

(iv) inexpensive

Ans. The electrolyte used is the solution of potassium argento cyanide, K[Ag(CN)₂], and anode is of pure silver. The following ions are formed in the solution.

K[Ag(CN)₂]₍aq₎ ⇌ K⁺₍aq₎ + [Ag(CN)₂]⁻₍aq₎

[Ag(CN)₂]⁻₍aq₎ ⇌ Ag⁺₍aq₎ + 2CN⁻₍aq₎

When the electric current is passed, Ag⁺ ions migrate towards the cathode and deposit after picking up electrons. Cyanide ions move towards the anode, where they react with silver to form AgCN, which combines with K(CN) to reproduce the electrolyte, K[Ag(CN)₂].

Ans. In copper plating, the anode is made of copper metal and the electrolyte is CuSO₄ to which a few drops of sulphuric acid are added to prevent its hydrolysis. The article to be electroplated is made the cathode. The following electrolytic reactions take place, at the electrodes.

At anode
Cu₍s₎ → Cu²⁺₍aq₎ + 2e⁻

At cathode
Cu²⁺₍aq₎ + 2e⁻ → Cu₍s₎

Ans. The cell which converts chemical energy into electrical energy is called galvanic cell.

Example Daniel Cell

Ans. An oxidation reaction takes place at zinc electrode and it is called anode while reduction takes place at copper electrode and it is called cathode.

Chemical reactions The following reactions take place at the two electrodes.

At anode
Zn₍s₎ → Zn²⁺₍aq₎ + 2e⁻ (Oxidation)

At cathode
Cu²⁺₍aq₎ + 2e⁻ → Cu₍s₎ (Reduction)

Overall cell reaction
Zn₍s₎ + Cu²⁺₍aq₎ → Zn²⁺₍aq₎ + Cu₍s₎

Ans. According to the available data the reactivity of metals depends upon how readily it is oxidized or lose electrons to form positive ions. If a galvanic cell is established between two metals, the more reactive metal will lose electrons and form positive ions.

Ans. If a cell is established between magnesium and copper it produces a higher voltage than a cell established between iron and copper. This will mean that magnesium is more reactive than iron because it has a tendency to lose electrons more easily.

Ans. Solution

Oxidation number (O.N) of O = -2
Oxidation number of Ca = +2
Oxidation number of Cℓ = x =?

O.N of Ca + 2(O.Nof Cℓ) + 6(O.Nof O) = 0
+2 + 2x + 6(-2) = 0
+2 + 2x - 12 = 0
2x = +12 - 2
2x = +10
x = 10/2 = +5

O.Nof Cℓ = x = +5

Ans. An oxidizing agent is a substance that gives oxygen or accepts electron from another substance. H₂O₂ Behaves as an oxidizing agent because it provides oxygen to other and accepts electron from other compound. In hydrogen peroxide, oxygen has an oxidation state of -1, which is unstable it readily gets reduced to -2 (as in water) causing the other substance to be oxidized.

Example

When it reacts with potassium iodide (KI) in acidic medium:

H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O

Iodide ions (I⁻) are oxidized to iodine (I₂) and H₂O₂ acts as an oxidizing agent.

Ans. Oxygen gas is produced: The difference in products is due to the fact that ions that get discharged at the electrodes depend on their concentration and ease of discharged. In both dilute and concentrated sodium chloride solutions, these species are present.

• Na⁺
• Cℓ⁻
• H⁺ and OH⁻ (from water)

At cathode (same in both cases)

Sodium is higher in electrochemical series, and being highly reactive nature of sodium water is reduced instead of sodium.

2H₂O + 2e⁻ → H₂ + 2OH⁻

At anode

In dilute solution Cℓ⁻ concentration is low. OH⁻ are preferentially discharged.

4OH⁻ → O₂ + 2H₂O + 4e⁻

In concentrated solution Cℓ⁻ ion concentration is high so Cℓ⁻ ions are discharged more easily than OH⁻

2Cℓ⁻ → Cℓ₂ + 2e⁻

Chlorine gas is produced.

So concentration affects which ions are preferentially discharged.

Ans. Corrosion becomes fast during rainy season because there is more moisture everywhere. This moisture creates a thin form of water and acts like an electrolyte letting electrons move easily to accelerate corrosion. Rain water contains more dissolved oxygen which also fuels oxidation (rusting in iron and steel).

Rain dissolves pollutants like SO₂, CO₂ and NOₓ from the air making acid rain. These acids increase corrosion rates especially in industrial areas.

Ans. In a Cu - Ag galvanic cell, silver is noble metal and has higher reduction potential so it gets reduced. So

At anode
Cu → Cu²⁺ + 2e⁻

Oxidation - Cu - electrode

At cathode
2Ag⁺ + 2e⁻ → Ag

Reduction - Ag - electrode

Ans. In a galvanic cell, the porous partition acts as a separator that allows ions to pass but prevents mixing of the solution. It is also called salt bridge. It performs the following functions.

• It mains electrical neutrality, by allowing ions to migrate between the two half - cells so charge does not build up and stop the reaction.
• It completes the internal circuit.
• Prevents direct mixing of electrolytes.
• Reduces diffusion of solutions.

Ans. See Q.3 from theory.

Ans.

A𝓁(NO₃)₃, H₂S₂O₇, Na₂S₂O₃

Ans. (i) A𝓁(NO₃)₃

O.N of A𝓁 = +3
O.N of O = -2
O.N of N = x

O.N of A𝓁 + 3(O.N of N) + 9(O.N of O) = 0
+3 + 3x + 9(-2) = 0
+3 + 3x - 18 = 0
3x - 15 = 0
3x = 15
x = 15/3 = +5

So O.N of N = +5

(ii) H₂S₂O₇

O.N of H = +1
O.N of O = -2
O.N of S = ? = x

2(+1) + 2(x) + 7(-2) = 0
+2 + 2x - 14 = 0
2x - 12 = 0
2x = 12

x = 12/2 = 6

So O.N of S = +6

(iii) Na₂S₂O₃

O.N of Na = +1
O.N of O = -2
O.N of S = x

2(+1) + 2(S) + 3(-2) = 0
+2 + 2x - 6 = 0
2x - 4 = 0
2x = 4

x = 4/2 = +2

So O.N of S = +2

Ans. Given the formulas of ionic compounds using ionic charges.

(i) Sodium hypochlorite

NaCℓO Sodium hypochlorite

(ii) Calcium oxalate

CaC₂O₄ Calcium oxalate

(iii) Ammonium phosphate

(NH₄)₃ PO₄ Ammonium phosphate

(iv) Calcium nitrite

Ca(NO₂)₂ Calcium nitrite

Ans. See Q.14 from theory.

Ans. See Q.10 from theory.

Ans. See Q.19 from theory.

Ans. See Q.16 from theory.