Understanding Redox Reactions: Oxidation and Reduction

• Redox reactions involve both reduction and oxidation.

Oxidation vs. Reduction

• Oxidation involves adding oxygen or removing hydrogen.
• Oxidation leads to the loss of electrons.
• Oxidation causes an increase in the oxidation number.
• Reduction involves removing oxygen or adding hydrogen.
• Reduction involves gaining electrons.
• Reduction causes a decrease oxidation number.

Oxidizing vs. Reducing Agents

• An oxidizing agent (oxidant) oxidizes another substance and is itself reduced.
• A reducing agent (reductant) reduces another substance and is itself oxidized.
• Examples of oxidizing agents include acidified KMnO4, acidified K2Cr2O7, chlorine, oxygen, and hydrogen peroxide.
• Examples of reducing agents include hydrogen, carbon, carbon monoxide, potassium iodide, and reactive metals.

Redox Reaction Examples

• Reduction of Hematite (Iron oxide to Iron):
  • Fe2O3 + 3CO → 2Fe + 3CO2
  • Fe2O3 is reduced (oxidizing agent).
  • CO is oxidized (reducing agent).
• Reaction between Chlorine and Hydrogen Sulfide:
  • Cl2 + H2S → 2HCl + S
  • Cl2 is reduced (oxidizing agent).
  • H2S is oxidized (reducing agent).
• Combustion of Magnesium:
  • 2Mg + O2 → 2MgO
  • Mg is oxidized (reducing agent): Mg → Mg2+ + 2e-
  • O2 is reduced (oxidizing agent): O2 + 2e- → O2-

OIL RIG mnemonic

• Oxidation Is Loss (of electrons)
• Reduction Is Gain (of electrons)

Oxidation Numbers

• The oxidation number indicates the degree of oxidation of an atom in a compound.
• Oxidation numbers can be positive, negative, or zero.
• An uncombined element has an oxidation number of zero.
• A monoatomic ion's oxidation number is equal to its charge.
• Group I elements in a compound always have an oxidation number of +1.
• Group II elements in a compound always have an oxidation number of +2.
• Fluorine in a compound always has an oxidation number of -1.
• Oxygen typically has an oxidation number of -2, except in peroxides (-1) and F2O (+2).
• Hydrogen typically has an oxidation number of +1, except in metal hydrides (e.g., NaH), where it is -1.
• The more electronegative element in a substance is assigned the negative oxidation number.
• The sum of oxidation numbers in a neutral compound is zero.
• The sum of oxidation numbers in an ion equals the charge of the ion.
• Metals generally form positive ions by losing electrons, increasing their oxidation number.
• Non-metals generally form negative ions by gaining electrons, decreasing their oxidation number.

Deducing Oxidation Numbers

To deduce oxidation numbers, use the known oxidation numbers of common elements (e.g., O, H, Group 1 metals) and the rules for the sum of oxidation numbers in a compound or ion.

Classifying Reactions Using Oxidation Numbers

• Compare the oxidation numbers of elements before and after a reaction.
• If the oxidation number of an element increases, it has been oxidized.
• If the oxidation number of an element decreases, it has been reduced.
• Redox reactions involve both oxidation and reduction.

Disproportionation Reactions

• Disproportionation reactions involve a single element being simultaneously oxidized and reduced.

Constructing Equations Using Oxidation Numbers

• Half-ionic equations represent the oxidation or reduction process, showing electron gain or loss.
• Full ionic equations are constructed by combining two half-equations, ensuring the electrons cancel out.

Balancing Complex Half-Ionic Equations

• Write the unbalanced equation for the species undergoing oxidation or reduction.
• Balance the atoms (other than O and H) that are being oxidized or reduced.
• Add H2O to balance oxygen atoms.
• Add H+ to balance hydrogen atoms.
• Add e- to balance the charge.

Balancing Redox Equations Using Oxidation Numbers

• Identify the elements whose oxidation numbers change.
• Determine the ratio of change in oxidation numbers.
• Balance the elements that change oxidation numbers.
• Balance the remaining elements, including H and O atoms.
• Check that the equation is balanced for both atoms and charge.

Group 1 & 2 Elements: Ionization Energies

• Ionization energy depends on electronic structure and affects physical/chemical properties.
• Ionization requires energy to overcome electrostatic attraction between electrons and protons.
• First ionization energy is the energy to remove one mole of electrons from one mole of gaseous atoms to form 1+ ions.
• Second ionization energy is the energy to remove one mole of electrons from one mole of gaseous +1 ions to form +2 ions.
• Factors affecting ionization energy: effective nuclear charge, atomic radius, and shielding by inner electrons. First ionization energy decreases down Groups 1 and 2 because atomic radius increases and shielding increases.

Trends in Reactivity Down Groups 1 and 2

• Reactivity increases down Group 1 and 2 due to the decreasing energy needed to remove electrons.
• Group I metals tarnish in air, forming a dull oxide layer.
• Products are oxides containing M+ and O2- ions.
• Lithium forms lithium oxide, sodium forms sodium oxide and peroxide.
• Sodium has a larger ionic radius than Lithium.
• Group 2 metals burn to form metal oxides when heated.
• Without heating, a slow reaction with oxygen forms a metal oxide coating.
• Products are oxides containing M2+ and O2- ions.

Reaction with Chlorine

• Group 1 and 2 metals react with chlorine (when heated) to form chlorides.
• The reaction becomes more vigorous down the groups.
• Reaction of Group 1 metals produces chlorides containing M+ and Cl- ions.
• Reaction of Group 2 metals produces chlorides containing M2+ and Cl- ions.

Reaction with Water

• Group I metals react vigorously with water.
• Products are hydrogen gas and metal hydroxides.
• The general formual is 2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)
• Lithium reacts immediately, forming lithium hydroxide and hydrogen gas.
• Reactions become more vigorous down the group
• Sodium sometimes produces a yellow flame due to ignition of hydrogen gas.
• Potassium usually causes hydrogen to catch fire.
• Group 2 metals react with water, with increasing vigor down the group (Mg slowly, Ca/Sr/Ba more).
• Products are hydrogen gas plus metal hydroxides, general formualis M(s) + 2H2O(l) → M(OH)2(aq) + H2(g.

Magnesium and Steam

• Magnesium reacts rapidly with steam when heated, creating magnesium oxide and hydrogen gas.
• Hydrogen formed is potentially dangerous.

Reactions of Oxides & Hydroxides of Group 1 & 2

• Group I and 2 oxides are basic oxides, resulting in alkaline solutions.
• Products form colourless solutions.
• O2- + H2O→2OH-

Reaction with Acids

• Group 1 & 2 oxides (hydroxides) react with acids to form a salt and water via neutralization.
• Reactions are exothermic & forms colourless solutions.

Uses of Lime

• In farming: Used to reduce soil acidity and increase crop yield.
• Neutralizes soil acidity
• Testing for carbon dioxide: used for testing Carbon dioxide.
• Suspension of magnesium hydroxide in water acts as antacid.

Group 2 Hydroxides

• In water it increases down the group, resulting in more alkaline solutions.

Group 2 Sulfates

• In water it decreases down the group
• Magnesium Sulfate is soluble
• Calcium Sulfate is slightly soluble
• Strontium Sulfate & Barium Sulfate are insoluble

Barium Compounds

• Solutions containing Barium test for presence of Sulfate ions in solutions.
• It is used in hospitals to better see tissues during X-rays.

Thermal Stability

• It is the measure of how much a compound decomposes when heated.
  • Very Thermally stable →doesn't decompose at all.
  • Not thermally stable → decompose as much as possible
• The charge and size of the Cation present influence the stability of Nitrate and Carbonate ions.
• A small positive ions will more readily polarize a carbonate ion vs large ions.

Thermal Decomposition of Carbonates

• All Group 1 and 2 carbonates are white solids, that when heated, they either don't decompose or decompose giving metal oxides and carbon dioxide gas.

Group

• Only Lithium Carbonate is the only Group I carbonate that decomposes to form Lithium oxide and Cabon dioxcied.
• Others will remin stable.

Group 2

• Will decompose heated giving off the metal oxide and carbon dioxide.
• Decomposition gets harder down the group.

Thermal decomposition of Group 1 & 2 Nitrates

• They are all white solids, that when heated, they either decompose to give: metal oxide + Nitrogen dioxide + Oxygen or decompose partially to give: Metal nitrite + Oxygen
• Metal Nitrates should be carried out in a fume cupboard because Nitrogen dioxide is a toxic gas

Group I

• It the Nitrate is only Group I nitrate that decomposes to produce Nitrogen dioxide. It breaks down into Lithium oxide, Nitrogen dioxide and Oxygen.
• The rest partially to give Metal nitrite + Oxygen

Testing for Oxygen & Nitrogen dioxide Gas produced

• Will relight a glowing splint.
• Nitrogen dioxide: it's a toxic brown gas, that when dissolved in water, it will give acidic solution.

Flame Tests

• Metal ions produce a colour if strongly heated in a flame
• It is used to identify metals in different solutions.

How to use

1. Dip a platinum loop or nichrome wire in hydrochloric acid solution then rinse it with water
2. Dip the clean loop in the sample solution.
3. Hold loop the the flame and look for color

Colors Produced

• red: Lithium
• yellow/range: Sodium
• lilac: Potassium
• red/purple: Rubidium
• blue/voilet : Caesium
• (no colour): Beryllium
• (no colour): Magnesium
• (brick) red: Calcium
• (crimson) red: Strontium
• Barium: (apple) green

Testing for Ions

• use drops of sodium hydroxide solution and heat
• It will release Pungent Ammonia gas, which turns litmus paper blue in the presencde of Ammoniun ions.

Carbonate/Hydrogencarbonate ions

• Adding dilutred Hydrochloric acid produces Carbon Dioxide gas, which when led throught lime water makes the water appear milky.

Titration

• Titration is used to determine unknown concentration of a solution.

Experimental Method

1. Rinse the conical flask with deionised water and place it on a white tile.
2. Using a pipette filler, rinse the pipette with deionised water and then with some of the hydrochloric acid solution.
3. Use the pipette to transfer 25 cm³ of HCl to the conical flask.
4. Add 3 drops of an indicator to the flask.
5. Rinse the burette with deionised water then with some of the sodium hydroxide.
6. Fill the burette with NaOH solution and set it up to stand above the conical flask.
7. Record the burette reading.
8. Add NaOH from the Burette slowly to the flask & swirl until the colour of indicator just changes.
9. Record the burette reading again.
10. Empty and rinse the flask with deionised water and repeat the titration until concordant titres have been obtained. Dr. Mar
11. Carry out calculations to find concentration of HCl. Com

Concordant titres

• It is considered a good expermintal technqiue to perform the titration multiple times to attain results with are similar with each other.

Indicators & Colours

• Methyl orange:
  • Red: Color in acid
  • Yellow: Color in alkali
• Phenol phthalein:
  • Colorless: Color in acid -Pink: Color in alkali

Important Tips on titration experiment

• Burette must be washed with water then with the titrant.
• Flask must be washed with water ONLY.

Measurement uncertainty

• It is the potential error when using a device.

Calculation of Measurement uncertainty

• when using a balance, weighing a sample of 5 g will lead to lower percentage uncertainty than weighing a 0.5 g sample.
• In a titration, having a titre value of 30 cm³ will have lower percentage uncertainty than titre value of 10 cm³

Group VII (Halogens

• Often referred to has Halogens. Comes from "salt producer" Most Recent F CL Br I At

Trend in Electronegativity

• Electronegativity is the ability of an atom to attract the bonding pairs of electrons in a covalent bond. Electrochemistry:
• Electronegativity of Group 7 is the highest of any group in the periodic table.
• Fluorine is the most electronegative of all elements.
• Electronegativity decreases down the group.

Halogen/ Halide displacement reaction

• A more reactive Halogen can displace a less reactive halogen from a Halide solution of the less reactive Halogen.
  • Chlorine displace Bromine & lodine
  • Bromine displaces lodine but not Chlorine
  • lodine can't displace either Chlorine or Bromine

Disproportionation reactions of Chlorine and Bromine and lodine

• reactions of Chlorine are reactions in which Chlorine undergoes both oxidation and reduction simultaneously
• React is similar ways with chlorine and with:
  • water
  • cold alkali
  • Hot alkali

Conc. Sulfuric

• Halogens act as oxidizing agents, and their oxidizing power decreases down the group.
• Halide ions act as reducing agents, and their reducing power increases down the group.
• Sulfuric acid can act as an oxidizing agent as well as an acid.
• The extent of it reduction depends on the following:
  • Sulfur
  • Sulfur dioxide
  • Hydrogen Sulde Halides with aq. Silver Nitrate •Aq. Silver nitrate solution is used to test for presence of Halide ions. •Dil. nitric acid is added first to the solution before adding silver nitrate to remove any other ions that may also form precipitates (specially carbonate ions.) •Solutions containing halide ions form precipitates when tested with the reagent aq. Silver nitrate Halides with water •All Hydrogen halide are colourless gases, and exist as polar diatomic molecules. Hydrogen halides react with water to form acidic solutions

Explore the fundamentals of redox reactions, including the definitions of oxidation and reduction. Learn about oxidizing and reducing agents with examples such as acidified KMnO4 and hydrogen. Understand reaction examples like the reduction of hematite.