Unit 1: Stoichiometry

Concept of relative atomic masses, molecular masses, and molecular formulae.

Calculation of relative atomic masses, relative molecular masses considering existence of isotopes, understanding empirical formula, molecular formula, mass and formula mass of different compounds.

Validation of law of stoichiometric proportions to introduce concept of moles. Includes mole-mole, mole-mass, mass-mass, mass-volume conversions.

Application of concept of moles to work out various stoichiometric proportions like % yield, % purity, reacting masses of reactants and products, limiting reagents, atom economy.

Extending the concept of mole to understand gaseous reactions and calculation of concentration of solutions through measuring volumes of solutions required by titration method, concentration numericals.

What is Stoichiometry?

Stoichiometry bridges the gap between macro quantities like mass and micro quantities like atoms and molecules.

Laws of Stoichiometry:

• Law of Conservation of Mass
• Law of Constant Proportions
• Law of Multiple Proportions
• Law of Reciprocal Proportions
• Law of Gaseous Volume

Moles

A mole is a fixed constant in chemistry. Similar to how one dozen equals twelve objects, one mole equals 6.02 × 10²³ atoms of an element or compound.

The atomic weight/relative atomic mass of an element is equal to its molar mass in grams. The molar mass of an element is the mass of 1 mole of it (6.02 × 10²³ atoms).

Example: The atomic weight/relative atomic mass, and thus molar mass in grams, of 1 mole of oxygen is 15.999 g/mol (~16.0 g/mol).

Two moles of oxygen = 16 × 2 = 32 g/mol.

Mole-Mole Relationships

In a balanced equation, the number of moles of a compound is given by the coefficient in front of it. By finding the ratio of the number of moles of reactants to products, it is possible to find the mass of products formed for a given mass of reactants.

Mole-Volume Relationships

The Bronsted-Lowry Theory

Here, acids are “proton donors”, meaning they give out H+ ions, which are basically just protons. A hydrogen atom has one proton and one electron with zero neutrons. Strip away the electron to get an H+ ion, and Bob’s your uncle, you have a proton.

Acids have an excess of H+ ions.

Meanwhile, bases are “proton receivers”, meaning they can accept H+ ions, OR they have an excess of OH- ions. Bases that are soluble in water are called alkali. Remember that not all bases are alkalis (that is to say, not all bases are soluble in water, but all alkalis are bases!)

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Conjugate Acid-Base Pair

Conjugate acid-base pairs are pairs with a single proton difference, or the difference of an H+ ion. Not two, not three, just one.

HF (hydrogen fluoride) is the acid here, because it is a proton donor and gives up a proton, causing it to become negatively charged fluorine (because bye bye hydrogen!). This makes the F- a conjugate base. F- now can accept a proton.

Meanwhile, H2O is the base here because it accepts the proton, causing it to become positively charged H3O (because hello hydrogen), making H3O a conjugate acid. H3O can now give up a proton.

The charges of the ions change because atoms can have their charges changed in two ways: either by an increase or decrease in electrons, or an increase or decrease in the number of protons. Here, it’s with protons!

Note that some chemical compounds can be both conjugate acids and bases in different conditions, like NH3

If we take it as a base:

If we take it as an acid:

Universal Indicator

All about Acids!

Acids can be weak or strong.

Strong Acids:

Dissociate completely, more acidic, lower pH, think HCl, H2SO4.

Weak Acids:

Dissociate partially, less acidic, higher pH, think acetic acid.

“Dissociation” refers to them breaking down into ions. Weak acids will mostly stay as molecules and not break down into ions. To figure out whether an acid is weak or strong, we can use several distinguishing tests, such as:

Tests to Distinguish Strong and Weak Acids

• Conductivity Test: Stronger acids are better conductors because they produce more ions due to complete dissociation.
• pH Test: Strong acids have a pH of 2 or less, while weak acids have a pH of 2–7.
• Rate of Reaction: Stronger acids are more highly reactive because they produce more ions and can react fully with bases.

Reactions of Acids

These are super important, so try to memorize them to save time in the exam.

• Acid + Base → Salt + Water
• Acid + Metal Hydroxide → Salt + Water
• Acid + Metal Oxide → Salt + Water
• Acid + Metal Carbonate (CO3) → Salt + Water + Carbon Dioxide (CO2)
• Acid + Metal Hydrogen Carbonate → Salt + Water + CO2
• Acid + Metal → Salt + Hydrogen Gas (produces “effervescence”)

Note: Water is only produced if there is a hydroxide ion, oxide, carbonate, or hydrogen carbonate present.

Example: Sulfuric Acid and Copper(II) Oxide

Break down the reactants into their constituent ions:

React the first cation with the second anion and the first anion with the second cation:

Write down the products and adjust coefficients as necessary:

To test for CO2, you can do the flame test (place a candle in gas; if it goes out, it’s CO2) or the limewater test (CO2 + Ca(OH)2 → CaCO3, turning the limewater milky).

pH Scale

pH stands for “Power of Hydrogen” or “Potential of Hydrogen”. It measures the hydrogen ion concentration per liter of solution:

For bases, pOH is the “Power of the Hydroxide Ion”:

The more diluted an acid, the higher the pH value. The more diluted a base, the lower the pOH value. Remember:

If pH increases by 1 unit, H+ concentration decreases by 10 units and vice-versa. The same applies for OH-.

Antacids

Used to neutralize excessive stomach acid and treat acid indigestion. Examples include:

• Milk of magnesia (Mg(OH)2, magnesium hydroxide)
• Alka-Seltzer (NaHCO3, sodium bicarbonate)

Nature of Oxides

There are 4 main types:

• Metallic oxides: Basic. React with acids to form salt + water. Example: MgO, Na2O. Dissolve in water to form alkalis (pH 11–14).
• Nonmetal oxides: Acidic. Dissolve in water to form acids (e.g., CO2 + H2O → H2CO3, SO2 + H2O → H2SO3). React with bases to form salt + water (e.g., 2NaOH + CO2 → Na2CO3 + 2H2O).
• Neutral oxides: Do not react with acids or bases (e.g., CO, N2O, H2O). Some can be dangerous, like CO reacting with hemoglobin.
• Amphoteric oxides: React with both acids and bases. Insoluble in water. Examples:

Strength Versus Concentration

The strength of an acid or base depends on how much it dissociates in water according to both the Arrhenius theory and the Bronsted-Lowry theory. Strong acids like HCl will completely dissolve in water to form H+ and Cl- ions, while weak acids like CH3COOH will only partially dissociate into CH3COO- ions and H+ ions.

Dilution

Always add acid to water in small volumes with continuous stirring so that water can gradually absorb the heat from the acid. The stirring evenly distributes the energy from the water molecules reacting with the acid. If you add acid all at once then the water will immediately turn into steam, damaging the beaker.

NEVER add water to acid because the reaction produces steam immediately, which can damage the beaker or cause injury. Adding acid to water is safer because the water can absorb more heat.

Dilution is calculated using the formula:

To calculate the volume of water to add, subtract V1 from V2.

Properties of Bases

• Base + Amphoteric Metal → Salt + Hydrogen
• Base + Nonmetal Oxide → Salt + Water
• Base + Ammonium Salts → Salt + Ammonia (NH3) + Water

Metal oxide bases are soluble in water (e.g., KOH). Transition metal oxide bases are insoluble in water (e.g., CuO).

Salts

Salts are ionic compounds made of positive and negative ions. They are formed when H+ from an acid is replaced by a metal ion from a base or metal. Salts can be soluble (aq) or insoluble (s).

Applications include fertilizers, batteries, fungicides, and health products:

• Potassium chloride (fertilizer, part of NPK)
• Magnesium sulphate (fungicide, epsom salt)
• Calcium carbonate (antacid)

Making Salts with Metals

React metal with acid to produce salt and hydrogen gas:

Making Salts with Insoluble Bases

React insoluble base with acid to produce salt and water. Example with copper oxide:

1. Heat diluted acid and stir continuously.
2. Add base to acid until solution saturates and precipitate forms.
3. Filter solution to remove excess base.
4. Evaporate solution to leave behind salt (CuSO4).

Making Salts with Soluble Bases (Alkalis) via Neutralization

React alkali with acid using titration to produce salt and water. Example:

Making Insoluble Salts Through Precipitation

Salts can be soluble or insoluble depending on composition.

Soluble	Insoluble
All sodium (NaCl), potassium (KOH), ammonium (NH4NO3) salts	All nitrates (NaNO3)
Chlorides (KCl)…	…except AgCl, PbCl2
Sulphates (Cu2SO4)…	…except CaSO4, BaSO4, PbSO4
Sodium (Na2CO3), potassium (K2CO3), ammonium ((NH4)2CO3) carbonates	…otherwise, all carbonates are insoluble

Example: Precipitation Reaction for Barium Sulphate

1. Mix BaCl2 (aq) and MgSO4 (aq).
2. Double displacement reaction forms precipitate (BaSO4) and soluble product (MgCl2).
3. Filter to remove BaSO4.
4. Wash precipitate to remove soluble impurities.
5. Dry precipitate in warm oven.

Criterion B: Reaction of Metals With Acids

Example: Zinc with HCl, H2SO4, CH3COOH (undiluted)

• IV: Strength of acid (HCl > H2SO4 > CH3COOH)
• DV: Rate of reaction measured by disappearance of zinc
• CV: Mass of zinc, volume of acid

Procedure

1. Gather materials:
   • Acetic acid (CH3COOH) 300 cm3
   • Hydrochloric acid (HCl) 300 cm3
   • Sulfuric acid (H2SO4) 300 cm3
   • Magnesium 10 g chunk
   • Beakers 300 ml (3)
   • Gloves, tweezers, stopwatch
2. Wear gloves and use tweezers for safety.
3. Add 100 cm3 of each acid to separate beakers.
4. Add 10 g magnesium to each beaker and measure time for complete reaction using a stopwatch.
5. Repeat for all acids, three trials per acid, calculate average time.

The experiment can be repeated with copper and dilute acids to compare reactions.

Crit B + C: Catalysts and H₂O₂

Variables

• Independent Variable: Type of catalyst used (MnO₂, Fe₂O₃, CuO, ZnO)
• Dependent Variable: Rate of decomposition of H₂O₂ OR volume of foam formed over time
• Control Variables:
  • Mass of catalyst used
  • Volume of H₂O₂ added
  • Concentration of H₂O₂

Research Question

How does the type of catalyst added to identical volumes of H₂O₂ affect the rate of decomposition as measured by the volume of foam formed over time?

Hypothesis

If the type of catalyst used is changed, then the rate of decomposition of H₂O₂ as measured by the volume of foam produced will change, because the oxidation state of the metal in the catalyst differs.

Manipulation of Variables

• IV: Type of catalyst used, manipulated using catalysts with increasing oxidation states (ZnO & CuO → Fe₂O₃ → MnO₂)
• DV: Rate of decomposition of H₂O₂ measured by foam volume
• CV1: Mass of catalyst kept constant
• CV2: Volume of H₂O₂ kept constant

Demonstration Video

Procedure

1. Gather materials:
   • 20 ml detergent
   • Pipette
   • 100 ml measuring cylinders (5)
   • 6% H₂O₂ (100 ml)
   • Manganese (IV) oxide, Zinc oxide, Copper (II) oxide, Iron (III) oxide (10 g each)
   • Gloves
2. Safety:
   • Wear gloves to avoid skin contact with H₂O₂
   • Keep measuring cylinder away from face to avoid fumes
3. Add 10 g of each catalyst to separate measuring cylinders.
4. Observe and record the volume of foam produced over time.

Observation Table

Conclusion

Extension

Real-Life Applications of Acids and Bases

Acids

Acid Rain: Oxides of sulfur and nitrogen dissolve in rainwater:

Sources: burning fossil fuels containing sulfur, e.g., at power stations.

Effects: damages crops, makes soil acidic, corrodes buildings (e.g., limestone).

Bases

Neutralize acidic soil using CaCO₃ (limestone), CaO (quicklime), Ca(OH)₂ (slaked lime/limewater). CaCO₃ is preferred as it is insoluble and remains in soil to neutralize acid.

Sink cleaning: Ammonia is used as a drain cleaner.

Experimental Chemistry

Acid can be added to a round-bottom flask using:

• Thistle funnel: extends to bottom of flask, allows acid to flow safely
• Dropping funnel: short tube with knob to adjust release speed, prevents bumping

Gas Collection and Testing

>A thistle funnel versus a dropping funnel setup.

To remove gas produced in a reaction, an exhaust pipe can be inserted into the stopper. There are 4 main methods:

• Upward air displacement (Downward delivery): Collects heavier gases that sink below lighter gases (e.g., HCl, SO₂).
• Downward air displacement (Upward delivery): Collects lighter gases that rise above heavier gases (e.g., H₂, NH₃).
• Over water: For gases insoluble in water (e.g., CO₂, O₂). Gas passes through water and collects in an upturned jar.
• Gas syringe: Measures the volume of any gas. Gas pressure pushes the plunger to measure volume produced.

Testing for Gases

• Ammonia (NH₃): Colourless, pungent, basic. Turns red litmus paper blue as NH₃ accepts H⁺ to form NH₄⁺.
• Carbon dioxide (CO₂): Colourless, odorless, mildly acidic. Turns limewater milky (CaCO₃ precipitate).
• Chlorine (Cl₂): Green, toxic. Bleaches moist pH indicator paper white.
• Hydrogen (H₂): Colourless, odorless, explosive with oxygen. Test: lit splint → "pop" sound.
• Oxygen (O₂): Colourless, odorless. Glowing splint reignites in pure O₂.

Ion Tests

Used to detect cations (+ve ions) or anions (-ve ions) in unknown salts through gas or precipitate formation.

Test for Cations

Cation	Test	Observation/Signs	Equation
Ammonium (NH4+)	Dilute NaOH + heat	Ammonia gas released, red litmus → blue	NH4+ + OH- → NH3 + H2O
Copper (II) (Cu2+)	Dilute NaOH or NH3	Pale blue ppt; dissolves with excess NH3 → dark blue solution	Cu2+ + 2OH- → Cu(OH)2 [Cu(NH3)4]2+ with excess NH3
Iron (II) (Fe2+)	Dilute NaOH or NH3	Pale green ppt	Fe2+ + 2OH- → Fe(OH)2
Iron (III) (Fe3+)	Dilute NaOH or NH3	Red-brown ppt	Fe3+ + 3OH- → Fe(OH)3
Aluminum (Al3+)	Dilute NaOH or NH3	White ppt; dissolves in excess NaOH → colorless	Al3+ + 3OH- → Al(OH)3 Al(OH)3 + OH- → [Al(OH)4]-
Zinc (Zn2+)	Dilute NaOH or NH3	White ppt; dissolves in excess → colorless	Zn2+ + 2OH- → Zn(OH)2 Excess → [Zn(OH)4]2- or [Zn(NH3)4]2+
Calcium (Ca2+)	Dilute NaOH or NH3	White ppt with NaOH; little/no ppt with NH3	Ca2+ + 2OH- → Ca(OH)2

Test for Anions

• Halides: Add dilute HNO₃ + AgNO₃. Chloride → white ppt, Bromide → cream ppt, Iodide → yellow ppt.
• Sulfates: Add dilute HCl + Ba(NO₃)₂. White ppt of BaSO₄ indicates sulfate.
• Nitrates: Add NaOH + Al foil + heat. Ammonia gas released confirms nitrate presence.
• Carbonates: Add dilute HCl; bubble gas through limewater. Milky solution confirms CO₂ and carbonate presence.

Types of Water

Water hardness measures dissolved minerals, mainly calcium and magnesium ions, in mg/L or ppm.

• Hard water: Contains calcium + magnesium salts (bicarbonates, chlorides, sulphates). Forms scum with soap, causes scale in pipes and kettles.
• Temporary hardness: Dissolved calcium/magnesium hydrogencarbonates. Removed by boiling.
• Permanent hardness: Dissolved calcium/magnesium sulphates. Cannot be removed by boiling.

Why water becomes hard: Flowing through chalk or lime deposits, water picks up calcium ions.

Unit 3: Organic Chemistry

Topics include:

• Hydrocarbons – saturated and unsaturated, homologous series, characteristics
• Alkanes & Alkenes – nomenclature, preparation, properties, reactions (combustion, substitution, addition, cracking)
• Alcohols & Carboxylic acids – preparation, properties, reactions (esterification)
• Polymers – natural (fats, oils, proteins, carbs) and synthetic (terylene, nylon)
• Soaps and detergents – saponification, hydrogenation of oils, issues with hard water
• New materials – molecular gastronomy, smart materials, nanoworld
• Environmental issues – greenhouse effect, global warming, smog, ozone hole, VOCs, paints & plastic pollution, leaded petrol

Carbon

Carbon is in group 4 of the periodic table and forms up to 4 bonds. Other groups:

• Group 1: 1 bond
• Group 2: 2 bonds
• Group 3: 3 bonds
• Group 5: 3 bonds
• Group 6: 2 bonds
• Group 7: 1 bond

Drawing Lewis Structures

Each line represents a shared pair of electrons. Every element wants 8 electrons in its outer shell (except H: 2 electrons).

Example: Fluoromethane (CH₃F). Carbon forms 4 bonds, each hydrogen 1, fluorine 1. Lone pairs shown on fluorine only.

Alkanes, Alkenes, and Alkynes

No. of Carbons	Prefix
1	Meth-
2	Eth-
3	Prop-
4	But-
5	Pent-
6	Hex-
7	Hept-
8	Oct-
9	Non-
10	Dec-

Alkanes

Saturated hydrocarbons (only single bonds). Formula: C_nH_2n+2. Example: 4 carbons → C₄H₁₀ = butane.

Physical properties:

• Insoluble in water (nonpolar)
• Soluble in nonpolar substances
• Floats on water (density <1 g/mL)

Chemical properties:

• Nonpolar
• Complete combustion → CO₂ + H₂O + heat
• Incomplete combustion → CO + soot

Alkenes & Alkynes

Unsaturated hydrocarbons with double/triple bonds. Alkenes formula: C_nH_2n. Example: 4 carbons → C₄H₈ = butene.

The formula only works if the alkene has a single double bond.

Alkenes Physical Properties

• Insoluble in water: Nonpolar, cannot dissolve in polar solvents like water.
• Soluble in nonpolar compounds: Dissolves nonpolar substances like oil.
• Floats on water: Density < water, explains oil spills.

Alkenes Chemical Properties

• Nonpolar: Mostly covalent bonds, low polarity between C and H.

Alkynes

General formula: C_nH_2n-2. Example: 4 carbons → C₄H₆ = butyne.

Rules apply only if there is exactly 1 triple bond. More triple bonds require different naming rules.

Naming Side Chains

Video reference: Click here

Step 1: Find the longest continuous carbon chain.

	Option #1: 3 carbons long
	Option #2: 4 carbons long
	Option #3: Also 4 carbons long
	Cat: 3 carbons long

The two longest chains are both 4 carbons → name = butane (single bonds).

Step 2: Number the chain forwards and backwards. Pick the least-numbered chain. CH₃ is attached to the 2nd carbon → methyl group.

Compound name: 2-methyl butane

Multiple Side Chains

Follow same rules: find the longest chain, number forwards/backwards, pick least numbers. Example with two CH₃ groups:

• Left to right: CH₃ at carbons 2 and 3
• Right to left: CH₃ also at 2 and 3
• Use smaller number first → 2,3-
• Two alkyl groups → prefix "di" → dimethyl

Final name: 2,3-dimethyl butane

Adding Halogens

Example compound:

Halogens in Organic Compounds

Instead of alkyl groups, halogens like chlorine can be attached. Use prefixes:

• Chlorine = chloro–
• Bromine = bromo–
• Fluorine = fluoro–
• Iodine = iodo–

Steps to name:

1. Find the longest carbon chain.
2. Number the chain forwards and backwards.
3. Pick the least-numbered chain.
4. Name the compound using halogen prefix + chain name.

Example with one chlorine: 2-chloro butane

Example with multiple halogens: alphabetical order → 1-bromo 2,3-dichloro butane

Naming Alkenes

Locate double bond, number to give lowest position. Example: double bond between carbons 2 and 3 → but-2-ene or 2-butene

Naming Alkynes

Locate triple bonds, number for lowest position. Two triple bonds → suffix –diyne. Example: triple bonds at carbons 1 and 3 → pent-1,3-diyne

Alcohols and Alkyls

Alkyls are groups with one less hydrogen than alkanes. Suffix: –yl

Alkane	Formula	Alkyl
Methane	CH4	CH3 (methyl)
Ethane	C2H6	C2H5 (ethyl)
Propane	C3H8	C3H7 (propyl)

Homologous Series

Members share the same functional group and differ by CH2 units.

• Alkanes, alkenes, alkynes, alcohols, carboxylic acids etc.
• Same chemical properties (functional group), different physical properties (mass, boiling/melting point).

Alkane	Alkene	Alkyne
CH4 (methane)	CH2 (methylene)	C2H2 (ethyne)
C2H6 (ethane)	C2H4 (ethylene)	C3H4 (propyne)
Difference: CH2	Difference: CH2	Difference: CH2

Fractional Distillation of Crude Oil

Video reference: Click here

Crude Oil

Crude oil is unprocessed oil, a mixture of hydrocarbons with varying chain lengths. Long-chain hydrocarbons have stronger London Dispersion forces and higher boiling points, making them harder to separate. Short-chain hydrocarbons have lower boiling points and are more valuable as fuels.

Fractional Distillation

Crude oil is heated into vapor and pumped into a fractionating column with a heat gradient (hot bottom, cooler top). Long-chain hydrocarbons condense lower in the column, while short-chain hydrocarbons condense higher. Bubble caps prevent mixing and encourage condensation. Each tray contains fractions with similar boiling points, used for different purposes like fuels or bitumen for roads.

Cracking

Long-chain hydrocarbons can be broken into smaller alkanes and alkenes. No fixed formula exists, but the number of atoms in the products equals the reactant.

• Thermal cracking: 750°C, 70 atm
• Catalytic cracking: 500°C, zeolite catalyst

Functional Groups

Functional groups define chemical properties of hydrocarbons:

• Alcohols
• Carboxylic acids
• Amides: Dicarboxylic Acid + Diamine → CONH linkage
• Esters: Dicarboxylic Acid + ...

Organic Polymers and Materials

• Kevlar: Polyamide stronger than steel; bulletproof vests, brake pads.
• Spandex / Lycra / Elastane: Stretchable polymer for clothing and bandages; light, UV-resistant, bacteria-resistant.
• Gore-Tex: Teflon sheet between fabrics; waterproof, breathable clothing.
• Thinsulate: Fine polypropylene fibers trap air for insulation, water-resistant.

Non-Polymers

Carbon fibres: Lightweight, strong material; used in sports equipment, bicycles, golf clubs, tennis rackets, gym equipment.