10th Class Chemistry Chapter 3 Notes

Definition

Oxidation number is the charge an atom would have if all its bonds were treated as fully ionic — assigned using a set of rules: an atom in its elemental form is $0$; a simple ion's oxidation number equals its charge; oxygen is usually $-2$ and hydrogen usually $+1$ (with known exceptions like peroxides); the oxidation numbers in a neutral compound must sum to $0$, and in a polyatomic ion must sum to the ion's charge.

Definition

Oxidation can be defined three equivalent ways: (1) loss of electrons, (2) gain of oxygen (or loss of hydrogen), (3) increase in oxidation number. Reduction is the reverse in each case: (1) gain of electrons, (2) loss of oxygen (or gain of hydrogen), (3) decrease in oxidation number.

$H_2{O}_2$ can act as an oxidizing agent because the oxygen in it (oxidation number $-1$, intermediate between $0$ and $-2$) can be reduced further to $-2$, gaining electrons from — and so oxidizing — whatever it reacts with.

Definition

An electrolytic cell uses electrical energy (from an external source) to force a non-spontaneous chemical reaction to occur — this is electrolysis. A galvanic (voltaic) cell does the opposite: it uses a spontaneous chemical reaction to generate electrical energy.

In an electrolytic cell, an electrolyte (a substance that conducts electricity when molten or dissolved) conducts current through the movement of its free ions towards the oppositely charged electrode, where they gain or lose electrons — unlike a metal conductor, where current is carried by free electrons alone.

The cathode is the electrode where reduction occurs (gain of electrons); electrons flow out of the electrolytic cell's power source, through the external circuit, into the cathode, where they're picked up by positive ions being reduced.

Definition

Electrolysis of dilute $H_{2}S{O}_4$ with Pt electrodes: at the cathode, $H^{+}$ ions are reduced to hydrogen gas; at the anode, $O{H}^{-}$/water is oxidized to oxygen gas — effectively electrolysing water itself, since Pt is inert (unreactive) and doesn't take part.

Electrolysis of $CuS{O}_4$ solution with copper (reactive) electrodes: at the cathode, $C{u}^{2+}$ is reduced and deposited as copper metal; at the anode, the copper electrode itself is oxidized and dissolves into solution as $C{u}^{2+}$ — this is exactly the principle behind copper plating, since the anode continuously replenishes the $C{u}^{2+}$ that plates onto the cathode.

Electrolysis of dilute NaCl solution: at the cathode, $H^{+}$ (from water) is preferentially discharged over $N{a}^{+}$, giving hydrogen gas; at the anode, $O{H}^{-}$ (from water) is preferentially discharged over $C{l}^{-}$ when dilute, giving oxygen gas.

Electrolysis of concentrated NaCl solution (brine): at the anode, the much higher concentration of $C{l}^{-}$ means chlorine gas is now discharged instead of oxygen; the cathode still produces hydrogen gas. This is exactly why dilute and concentrated NaCl solutions give different anode products — ion concentration, not just identity, affects which ion is preferentially discharged.

Electrolysis of dilute $ZnC{l}_2$ solution: at the cathode, $H^{+}$ (from water) is discharged in preference to $Z{n}^{2+}$, giving hydrogen gas, because a metal ion is only preferentially discharged over hydrogen at the cathode if that metal lies below hydrogen in the electrochemical series — zinc lies above hydrogen, so hydrogen wins.

Definition

Electrolysis of molten lead(II) chloride ($PbC{l}_2$): must be heated above its melting point ($\approx 501^{\circ }$C) so its ions are free to move (a solid's ions are fixed in place and can't conduct). At the cathode: $P{b}^{2+}+2e^{-}\to Pb$ (lead metal deposited). At the anode: $2C{l}^{-}\to C{l}_2+2e^{-}$ (chlorine gas released).

Definition

The electrochemical series ranks metals by how easily they lose electrons (get oxidized) — a metal higher in the series is more reactive and more readily gives up electrons than one lower down. This directly predicts displacement reactions: a higher (more reactive) metal will displace a lower (less reactive) metal's ions from solution.

Definition

A Daniel cell (Zn–Cu galvanic cell) has a zinc electrode in $ZnS{O}_4$ solution and a copper electrode in $CuS{O}_4$ solution, connected by a wire (external circuit) and a porous partition (or salt bridge) between the two solutions.

The porous partition allows ions to slowly pass between the two half-cells to maintain electrical neutrality in each solution as the reaction proceeds, without letting the solutions bulk-mix — without it, charge would build up in each half-cell and current would quickly stop flowing.

Zinc (more reactive) is the anode, oxidized: $Zn\to Z{n}^{2+}+2e^{-}$. Copper (less reactive) is the cathode, reduced: $C{u}^{2+}+2e^{-}\to Cu$.

Definition

A standard Daniel cell produces a standard voltage of about $1.1$ V under standard conditions.

Definition

A hydrogen–oxygen fuel cell produces electricity directly from the chemical reaction of hydrogen and oxygen combining to form water, using an alkaline electrolyte (commonly potassium hydroxide, KOH solution) between the electrodes. At the cathode (with oxygen supplied), oxygen is reduced; at the anode (with hydrogen supplied), hydrogen is oxidized, and the overall reaction produces water as the only by-product.

Fuel cells are considered environmentally friendly because their only major by-product is water, unlike burning fossil fuels, which releases carbon dioxide and other pollutants.

Definition

Electroplating deposits a thin, even layer of one metal onto the surface of another object using electrolysis, mainly to improve appearance (a shiny finish), resist corrosion, or reduce cost (coating a cheaper base metal with a thin layer of a more expensive one, like silver or gold plating).

A suitable electroplating electrolyte should contain ions of the plating metal, be soluble/conducting, and not react unwantedly with the electrodes or the object being plated.

In silver plating, the electrolyte is a silver salt solution (e.g. silver nitrate/silver cyanide complex) with a silver anode (which dissolves to keep replenishing $A{g}^{+}$ ions) and the object to be plated as the cathode. Copper plating follows the same principle: a copper salt electrolyte (e.g. copper sulfate) with a copper anode and the object as the cathode.

Definition

Corrosion (rusting, for iron) is the gradual oxidation of a metal by its environment. Iron rusts through reaction with both oxygen and water together — iron is oxidized to $F{e}^{2+}$, which is further oxidized to $F{e}^{3+}$, ultimately forming hydrated iron(III) oxide (rust).

Definition

Galvanization protects iron/steel by coating it with a layer of zinc — zinc is more reactive than iron, so even if the coating is scratched, the zinc corrodes preferentially and protects the exposed iron underneath (sacrificial protection).

Zinc (and sometimes magnesium) is the metal commonly used to provide sacrificial protection to buried steel pipelines and ship hulls, attached as blocks that corrode away in place of the structure they're protecting.