Chem Sem 1 Exam Revision

Who discovered the electron and how?

J.J. Thomson (1897) — cathode ray tube; beam deflected by electric/magnetic fields proving negative subatomic particles exist

Who discovered the nucleus and how?

Rutherford (1911) — gold foil experiment; most alpha particles passed through, ~1/8000 bounced back → small dense positive nucleus

What did Bohr contribute?

Electrons orbit nucleus in fixed energy shells; electrons emit/absorb specific energy as photons when changing shells — explains line spectra

What did Chadwick discover?

Neutron (1932) — bombarded beryllium with alpha particles; detected massive uncharged radiation

What was unknown at the time of Bohr's model (1913)?

Neutrons — discovered by Chadwick in 1932

Number of emission lines when electron excited to shell n

n(n−1)/2. e.g. excited to n=4: 4×3/2 = 6 lines

Which transition gives shortest wavelength emission?

Largest energy drop — e.g. n=4 → n=1 (more energy = shorter wavelength)

Which transition gives longest wavelength in hydrogen visible spectrum?

n=3 → n=2 (smallest energy gap in visible range; red line)

Energy gap trend between shells

Gets smaller as shells are further from nucleus

Ion notation format

Mass number top-left, atomic number bottom-left, charge top-right: e.g. ²³⁸₉₂U⁴⁺

For a cation, how many electrons?

Protons minus charge. e.g. K⁺ (Z=19): 19−1 = 18 electrons

For an anion, how many electrons?

Protons plus charge magnitude. e.g. O²⁻ (Z=8): 8+2 = 10 electrons

Electron configuration of K (Z=19)

2, 8, 8, 1

Electron configuration of Ca (Z=20)

2, 8, 8, 2

Electron configuration of Cl (Z=17)

2, 8, 7

Electron configuration of Ar (Z=18)

2, 8, 8

What are isotopes?

Same element (same protons), different neutrons → different mass number. Same chemical properties, different physical properties

Atomic radius trend across a period

Decreases left to right — same number of shells but increasing nuclear charge pulls electrons closer

Atomic radius trend down a group

Increases — extra shell added, valence electrons further from nucleus with more shielding

Melting point trend across Period 3

Na→Mg→Al increases (stronger metallic bonding); Si highest (covalent network); P, S, Cl, Ar drop sharply (covalent molecular — weak dispersion forces only)

Melting point trend down Group 1

Decreases — atomic radius increases, metallic bond weakens

Melting point trend down Group 17

Increases — larger molecules, more electrons, stronger dispersion forces between molecules

Melting point trend down Group 18

Increases — larger atoms, more electrons, stronger dispersion forces

Valency trend across Period 3

Na=1, Mg=2, Al=3, Si=4, P=3, S=2, Cl=1, Ar=0

Which Period 3 element has the highest melting point and why?

Si — covalent network solid; all covalent bonds throughout the lattice must be broken to melt it

Why do P, S, Cl, Ar have much lower melting points than Si?

They are covalent molecular — only weak dispersion forces between molecules need to be overcome, not strong covalent bonds

Which Group 2 element has the highest melting point?

Beryllium (Be) — smallest atomic radius, strongest metallic bonding

How to identify a metallic substance from properties

Conducts electricity in both solid AND liquid state; malleable; generally high melting point

How to identify an ionic substance from properties

Does NOT conduct as solid; DOES conduct when molten or dissolved; high melting point; hard and brittle

How to identify a covalent molecular substance from properties

Low melting point; does not conduct in solid or liquid state; exists as discrete molecules

How to identify a covalent network substance from properties

Very high melting point; extremely hard; does not conduct electricity (except graphite)

Metal + non-metal → what bond type?

Ionic

Non-metal + non-metal → what bond type?

Covalent

Metal alone or two metals → what bond type?

Metallic

Does SiO₂ conduct electricity?

No — covalent network solid; no mobile charged particles

Does graphite conduct electricity?

Yes — one delocalised electron per carbon atom free to move between layers

Ammonium ion

NH₄⁺

Hydroxide ion

OH⁻

Nitrate ion

NO₃⁻

Carbonate ion

CO₃²⁻

Sulfate ion

SO₄²⁻

Phosphate ion

PO₄³⁻

Permanganate ion

MnO₄⁻

Dichromate ion

Cr₂O₇²⁻

Chromate ion

CrO₄²⁻

Oxalate ion

C₂O₄²⁻

Cyanide ion

CN⁻

Hydrogen carbonate (bicarbonate) ion

HCO₃⁻

Acetate (ethanoate) ion

CH₃COO⁻

Thiosulfate ion

S₂O₃²⁻

Iron(II) and Iron(III) charges

Fe²⁺ and Fe³⁺

Copper(I) and Copper(II) charges

Cu⁺ and Cu²⁺

Lead(II) charge

Pb²⁺

Silver ion

Ag⁺

Zinc ion

Zn²⁺

Fe³⁺ colour in solution

Pale brown/yellow

Cu²⁺ colour in solution

Blue

Fe²⁺ colour in solution

Pale green

Acid + carbonate → products

Salt + water + CO₂(g)

Acid + metal → products

Salt + H₂(g)

Acid + base/metal hydroxide → products

Salt + water

Silver carbonate + nitric acid — net ionic equation

Ag₂CO₃(s) + 2H⁺(aq) → 2Ag⁺(aq) + H₂O(l) + CO₂(g)

Copper sulfate + excess KOH — observation

Blue precipitate of Cu(OH)₂ forms

Zinc + dilute H₂SO₄ — observation

Zinc dissolves; bubbles of hydrogen gas produced

Magnesium + hydrochloric acid — ionic equation

Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g)

Silver chloride precipitate colour

White

Silver bromide precipitate colour

Cream/pale yellow

Silver iodide precipitate colour

Yellow

Copper(II) hydroxide precipitate colour

Blue

Iron(III) hydroxide precipitate colour

Brown/rust

Iron(II) hydroxide precipitate colour

Green

Mole definition

6.022×10²³ particles (Avogadro's number)

n = m/M

moles = mass ÷ molar mass

n = V/22.71

moles of gas = volume in litres ÷ 22.71 at STP (0°C, 100 kPa)

PV = nRT — what is R?

8.314 J mol⁻¹ K⁻¹ (use P in Pa, V in m³, T in K)

How to convert °C to K

Add 273

Empirical formula method from % composition

Assume 100 g sample; convert % to g; divide each by molar mass to get moles; divide all by smallest value; round to whole number ratio

Molecular formula from empirical formula

Divide given molar mass by empirical formula mass; multiply all subscripts by that factor

Percentage composition formula

(mass of element in one mole ÷ molar mass of compound) × 100

General formula — alkanes

CₙH₂ₙ₊₂

General formula — alkenes

CₙH₂ₙ

General formula — cycloalkanes

CₙH₂ₙ

General formula — cycloalkenes

CₙH₂ₙ₋₂

What makes a hydrocarbon saturated?

Contains only single C–C bonds (alkanes, cycloalkanes)

What makes a hydrocarbon unsaturated?

Contains at least one C=C double bond (alkenes, cycloalkenes) or aromatic ring (benzene)

IUPAC naming — which end do you number from?

The end closest to the first branch or double bond

How to name a halogen substituent in IUPAC

Prefix: fluoro-, chloro-, bromo-, iodo- with position number; listed alphabetically before carbon chain name

Structural isomers definition

Same molecular formula, different structural arrangement of atoms

Geometric (cis/trans) isomers — when are they possible?

Only when each carbon of the C=C has two DIFFERENT groups attached

cis isomer definition

Same priority groups on the same side of the double bond

trans isomer definition

Same priority groups on opposite sides of the double bond

Alkene + Br₂ (no UV, no catalyst) — reaction type and observation

Addition reaction; bromine water/liquid bromine decolourises

Alkane + Cl₂ or Br₂ — conditions required

UV light (free radical substitution)

Benzene + halogen — reaction type and conditions

Electrophilic substitution; requires AlX₃ halogen carrier catalyst

Does benzene decolourise bromine water?

No — benzene does not undergo addition reactions; substitution only

Test to distinguish an alkene from a cycloalkane

Add bromine water — alkene decolourises it; cycloalkane does not react (no UV light present)

Cyclohexene + Br₂(l) no UV — product

1,2-dibromocyclohexane (addition across the double bond)

Benzene + I₂ + AlI₃ — organic product

Iodobenzene (substitution; HI also produced as byproduct)

Complete combustion of a hydrocarbon — products

CO₂(g) and H₂O(g) only (requires excess O₂)

Incomplete combustion of a hydrocarbon — products

CO and/or C(s) soot and H₂O (limited O₂)