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CAPE Chemistry > Period 3, Group ii, Group iV, Group Vii, Transition elements and Identification ofcations and anions. > Flashcards

Period 3, Group ii, Group iV, Group Vii, Transition elements and Identification ofcations and anions. Flashcards

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1
Q

What causes the variation in melting points across Period 3?

A

Na, Mg, Al: Giant metallic structures; melting points increase Na → Mg → Al because of increasing ionic charge and delocalized electrons.

Si: Giant covalent structure; very high melting point due to strong covalent bonds.

P, S, Cl: Simple molecular structures; low melting points due to weak van der Waals forces.

Ar: Exists as isolated atoms; very low melting point due to weak van der Waals forces.

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2
Q

What are the melting points and states of Period 3 elements at room temperature?

A

Na: 371 K, solid

Mg: 922 K, solid

Al: 933 K, solid

Si: 1683 K, solid

P: 317 K, solid

S: 392 K, solid

Cl: 172 K, gas

Ar: 84 K, gas

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3
Q

How does electrical conductivity vary across Period 3?

A

Na, Mg, Al: Good conductors (metallic bonding).

Si: Semiconductor (poor conductor at room temp).

P, S, Cl, Ar: Very poor conductors (simple molecular or atomic structures).

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4
Q

How does electronegativity change across Period 3?

A

It increases from Na to Cl:

Na: 0.9

Mg: 1.2

Al: 1.5

Si: 1.8

P: 2.1

S: 2.5

Cl: 3.0

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5
Q

How does atomic radius change across Period 3?

A

Atomic radius decreases across the period due to increasing nuclear charge without additional shielding.

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6
Q

How does ionic radius change across Period 3?

A

Metal cations (Na⁺ → Si⁴⁺): Ionic radius decreases across the period.

Non-metal anions (P³⁻ → Cl⁻): Ionic radius also decreases across the period.

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7
Q

Describe the reaction of Period 3 elements with water.

A

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) (vigorous)

Mg(s) + H₂O(g) → MgO(s) + H₂(g) (reacts with steam)

2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g) (steam, not cold/hot water)

Si, P, S: No reaction.

Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq) (slight reaction forming acids)

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8
Q

Describe the reaction of Period 3 elements with oxygen.

A

2Na(s) + O₂(g) → Na₂O(s) (or Na₂O₂)

2Mg(s) + O₂(g) → 2MgO(s)

4Al(s) + 3O₂(g) → 2Al₂O₃(s)

Si(s) + O₂(g) → SiO₂(s) (slow)

4P(s) + 5O₂(g) → P₄O₁₀(s)

S(s) + O₂(g) → SO₂(g)

Cl₂ and Ar: No reaction.

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9
Q

Describe the reaction of Period 3 elements with chlorine.

A

2Na(s) + Cl₂(g) → 2NaCl(s)

Mg(s) + Cl₂(g) → MgCl₂(s)

2Al(s) + 3Cl₂(g) → 2AlCl₃(s)

Si(s) + 2Cl₂(g) → SiCl₄(l)

2P(s) + 5Cl₂(g) → 2PCl₅(s)

S(s) + Cl₂(g) → S₂Cl₂(l)

Argon: No reaction.

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10
Q

How does oxidation state vary across Period 3 oxides and chlorides?

A

Oxides: Max oxidation state increases from +1 (Na₂O) to +7 (Cl₂O₇).

Chlorides: Max oxidation state rises from +1 (NaCl) to +5 (PCl₅); sulfur forms only up to +2.

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11
Q

What type of bonding is present in Period 3 oxides and chlorides?

A

Na₂O, MgO: Ionic

Al₂O₃: Ionic with covalent character

SiO₂: Giant covalent

P₄O₁₀, SO₂, Cl₂O: Simple molecular (covalent)

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12
Q

How do Period 3 oxides react with water?

A

Na₂O + H₂O → 2NaOH

MgO + H₂O → Mg(OH)₂ (slightly soluble)

Al₂O₃: No reaction with water (but reacts with acids/bases → amphoteric)

SiO₂: No reaction with water

P₄O₁₀ + 6H₂O → 4H₃PO₄

SO₂ + H₂O → H₂SO₃

SO₃ + H₂O → H₂SO₄

Cl₂O + H₂O → 2HClO

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13
Q

How does acid-base character change across Period 3 oxides?

A

Na₂O, MgO: Basic oxides

Al₂O₃: Amphoteric oxide

SiO₂, P₄O₁₀, SO₂, SO₃, Cl₂O: Acidic oxides

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14
Q

How do atomic and ionic radii change down Group II?

A

Both metallic and ionic radii increase down the group because of the addition of electron shells.

Ionic radii are smaller than metallic radii since outer shell electrons are lost when forming ions.

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15
Q

How do melting points vary down Group II?

A

High melting points (giant metallic structures).

General decrease down the group (except Mg, which has a different lattice structure).

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16
Q

How does density vary down Group II?

A

Density increases from Ca to Ba.

Be and Mg have higher densities because of efficient packing.

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17
Q

How do ionization energies change down Group II?

A

Decrease down the group because outer electrons are further from the nucleus and more shielded, making them easier to remove.

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18
Q

Describe the reaction of Group II metals with oxygen.

A

Equation example:
2Mg(s) + O₂(g) → 2MgO(s)

Reactivity with oxygen increases down the group.

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19
Q

Describe the reaction of Group II metals with water.

A

Mg: Reacts very slowly with cold water, more vigorously with steam.

Ca: Reacts vigorously with cold water.
Equation example:
Ca(s) + 2H₂O(l) → Ca(OH)₂(s) + H₂(g)

The reaction is generally exothermic, with the reactivity increasing as you move down the group. Magnesium reacts with steam to form magnesium oxide and hydrogen, while calcium, strontium, and barium react with water at room temperature. Beryllium is the only alkaline earth metal that does not react with water

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20
Q

Describe the reaction of Group II metals with dilute acids.

A

Example:
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Ba(s) + H₂SO₄(aq) → BaSO₄(s) + H₂(g)

Reactivity increases down the group.

decreasing ionization energy and increasing atomic radius

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21
Q

What happens to the solubility of Group II sulfates down the group?

A

Decreases down the group.

MgSO₄ is very soluble; BaSO₄ is insoluble.

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22
Q

What is a buffer solution?

A

A buffer solution is a mixture that resists changes in pH when a small amount of acid or base is added, or when it is diluted. It’s composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.

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23
Q

Why does sulfate solubility decrease down Group II?

A

Lattice energy decreases slightly; hydration enthalpy decreases more significantly.

ΔHsolution becomes more endothermic.

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24
Q

What is the test for sulfate ions using barium chloride?

A

Acidify solution with nitric or hydrochloric acid.

Add BaCl₂(aq).

A white precipitate of BaSO₄ forms if sulfate ions are present.
Equation:
Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

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25
Describe the thermal decomposition of Group II carbonates.
CaCO₃(s) → CaO(s) + CO₂(g) Resistance to decomposition increases down the group.
26
Describe the thermal decomposition of Group II nitrates.
Example: 2Mg(NO₃)₂(s) → 2MgO(s) + 4NO₂(g) + O₂(g) require higher temperatures to decompose. Resistance to decomposition increases down the group.
27
# thermal stability increases Name important uses of Group II compounds. ## Footnote The polarizing power of the Group II metal ions decreases down the group. This leads to less distortion of the nitrate ion, making the nitrates more stable.
Calcium oxide (CaO): Making cement and mortar; drying agent. Calcium hydroxide (Ca(OH)₂): Neutralizing acidic soils, bleaching powder, limewater. Calcium carbonate (CaCO₃): Building material, blast furnace slag removal. Magnesium oxide (MgO): Lining furnaces.
28
How does structure and bonding vary across Group IV?
C (diamond): Giant covalent, non-conductor. Si: Giant covalent, semi-conductor. Ge: Giant covalent, semi-conductor. Sn, Pb: Metallic, electrical conductors.
29
How do melting points change down Group IV?
High for C, Si, Ge (giant covalent). Decrease sharply for Sn, Pb (weaker metallic bonding due to larger ions).
30
How does electrical conductivity vary down Group IV?
C (diamond): Non-conductor. C (graphite): Conductor (delocalized electrons). Si, Ge: Semi-conductors. Sn, Pb: Good conductors (delocalized electrons).
31
Describe the bonding and structure of Group IV tetrachlorides (XCl₄).
Simple covalent molecules. Tetrahedral structure. Volatile liquids with low boiling points. Non-polar overall.
32
How do Group IV tetrachlorides react with water?
All hydrolyze except CCl₄. Produce oxides and HCl fumes. Example: SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g)
33
What trends exist in hydrolysis of Group IV tetrachlorides?
Ease of hydrolysis increases down the group (SiCl₄ → PbCl₄). Metallic character increases down the group.
34
What are the structures of Group IV oxides in +2 and +4 states?
Oxidation State | Structure | +2 | Simple molecular or ionic (GeO, SnO, PbO) | +4 | Giant covalent (SiO₂, GeO₂, SnO₂, PbO₂) except CO₂ (molecular gas) |
35
How does stability vary for Group IV oxides?
Stability of +4 oxides decreases down the group. Stability of +2 oxides increases down the group.
36
How does acid-base character change for Group IV oxides?
| Oxidation State | Trend | +2 | Amphoteric, more basic down the group (GeO < SnO < PbO). | +4 | Acidic → amphoteric (CO₂ acidic; PbO₂ amphoteric). |
37
Give examples of reactions of Group IV oxides with acids and alkalis.
SnO(s) + 2HCl(aq) → SnCl₂(aq) + H₂O(l) SnO(s) + 2OH⁻(aq) → SnO₂²⁻(aq) + H₂O(l)
38
How does thermal stability change for Group IV oxides?
+4 oxides become less thermally stable down the group. PbO₂ decomposes readily: PbO₂(s) → PbO(s) + ½O₂(g)
39
How does the relative stability of Group IV cations change?
+4 ions become less stable down the group. +2 ions become more stable down the group. Due to inert pair effect.
40
How can electrode potentials (E°) explain Group IV ion stability?
The more the value of standard reduction potential is negative, the more that element is stable. Pb⁴⁺ is a good oxidizing agent (accepts electrons). CO is a good reducing agent (donates electrons). +4 → +2 reduction becomes easier down the group.
41
# [](http://) What are some uses of ceramics based on silicon(IV) oxide (SiO₂)?
Furnace linings (thermal insulator, high melting point). Abrasives (hard, high melting point). Glass and porcelain (high melting point, moldable).
42
How do the physical properties of halogens vary down Group VII?
Volatility decreases (less volatile). Colour darkens: F₂: Pale yellow Cl₂: Yellow-green Br₂: Red-brown I₂: Grey-black (purple vapor). State at 20°C: F₂ and Cl₂: gases Br₂: liquid I₂: solid
43
How does atomic radius change down Group VII?
Increases down the group due to added electron shells.
44
How do melting and boiling points vary down Group VII?
Increase down the group due to stronger van der Waals forces (more electrons → more intermolecular forces).
45
How does density change down Group VII?
Increases down the group as atomic mass increases.
46
How does chemical reactivity vary for the halogens?
Reactivity as oxidizing agents decreases down the group. Cl₂ > Br₂ > I₂ as oxidizing agents.
47
Describe halogen displacement reactions.
A more reactive halogen displaces a less reactive halogen from solution: Example: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq) (orange color from Br₂)
48
How can E° values explain halogen reactivity?
Higher E° values = better oxidizing agents. E°(Cl₂/Cl⁻) > E°(Br₂/Br⁻) > E°(I₂/I⁻)
49
How does chlorine react with water?
Forms hydrochloric acid and chloric(I) acid: Cl₂(g) + H₂O(l) → H⁺(aq) + Cl⁻(aq) + ClO⁻(aq)
50
What is bromine water used to test for?
C=C double bonds (unsaturation) in organic compounds. Bromine water is primarily used to test for the presence of unsaturation in hydrocarbons, specifically alkenes Orange bromine water turns colourless.
51
Describe the reaction of halogens with hydrogen.
H₂ + X₂ → 2HX Reactivity decreases down the group: F₂ reacts explosively even in cold. Cl₂ reacts explosively in sunlight. Br₂ and I₂ react slowly with heat.
52
How does the thermal stability of hydrogen halides change down Group VII?
Decreases down the group (HF most stable, HI least stable). Due to decreasing bond energy.
53
What is the bond energy trend for hydrogen halides?
H-F: 562 kJ/mol H-Cl: 431 kJ/mol H-Br: 366 kJ/mol H-I: 299 kJ/mol
54
How do halide ions react with aqueous AgNO₃ and NH₃?
| Halide | AgNO₃ test | Reaction with NH₃ | Cl⁻ | White ppt (AgCl) | Dissolves in dilute NH₃ | | Br⁻ | Cream ppt (AgBr) | Dissolves in conc. NH₃ | | I⁻ | Yellow ppt (AgI) | Insoluble in NH₃
55
How do halide ions react with concentrated H₂SO₄?
Cl⁻: Only acid-base reaction → HCl gas. Br⁻: Reduces H₂SO₄ to SO₂ → Br₂ gas and SO₂ formed. I⁻: Reduces H₂SO₄ to SO₂, S, H₂S → I₂ formed.
56
Write the reaction of Cl⁻ with concentrated H₂SO₄.
NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g)
57
Write the redox reaction of Br⁻ with concentrated H₂SO₄
2HBr(g) + H₂SO₄(l) → SO₂(g) + Br₂(l) + 2H₂O(l)
58
Write the redox reaction of I⁻ with concentrated H₂SO₄.
Examples: 2HI(g) + H₂SO₄(l) → SO₂(g) + I₂(s) + 2H₂O(l) 8HI(g) + H₂SO₄(l) → H₂S(g) + 4I₂(s) + 4H₂O(l)
59
How does chlorine react with cold dilute NaOH?
Cl₂(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l) (Disproportionation reaction)
60
How does chlorine react with hot concentrated NaOH?
3Cl₂(aq) + 6NaOH(aq) → 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l) (Disproportionation reaction)
61
What is a transition element?
A d-block element that forms one or more stable ions with incomplete d-orbitals.
62
What are the key characteristics of transition elements?
Variable oxidation states Complex ion formation Colored compounds Catalytic activity Magnetic properties (paramagnetism/ferromagnetism)
63
Why zinc is not a tranitionm metal
While it is a member of the d-block and shares some properties with transition metals, it does not have an incompletely filled d-orbital in its common oxidation state, Zn2+.
64
How do transition metals compare with a typical s-block element (like calcium)?
| Property | Transition Metal | Calcium (s-block) | | Melting point | Higher | Lower | Density | Higher | Lower | Atomic radius | Smaller | Larger | Ionic radius | Smaller | Larger | First ionization energy | Higher | Lower | Electrical conductivity | Good | Good
65
How do you determine the electronic configurations of first row transition elements and their ions?
4s orbital is filled before 3d orbital. On ionization, 4s electrons are removed first. Example: Fe: [Ar] 3d⁶ 4s² Fe²⁺: [Ar] 3d⁶
66
Why do atomic and ionic radii, and ionization energies of transition elements change little across the period?
Shielding by d-electrons is poor. Increasing nuclear charge is balanced by slight contraction, resulting in small changes.
67
Why are many transition element ions colored?
d-orbital splitting in complexes causes certain wavelengths of light to be absorbed. Remaining light is transmitted/reflected, giving color.
68
What causes the color of transition metal complexes?
Light promotes an electron from a lower-energy d-orbital to a higher-energy d-orbital (d-d transition) in octahedral or other complexes.
69
How does oxidation state vary for vanadium?
Vanadium forms oxidation states: +2 (V²⁺) +3 (V³⁺) +4 (VO²⁺) +5 (VO₃⁻)
70
What are the colours of vanadium ions in different oxidation states?
V²⁺: Violet V³⁺: Green VO²⁺: Blue VO₃⁻: Yellow
71
What experimental method reduces vanadium(V) to lower oxidation states?
React acidified ammonium vanadate(V) solution with zinc metal. (Vanadium is reduced stepwise from +5 to +2.)
72
What shapes do transition element complexes commonly have?
Octahedral (6 ligands) Tetrahedral (4 ligands) Square planar (4 ligands, eg. Pt²⁺ complexes)
73
What are examples of octahedral, tetrahedral, and square planar complexes?
Octahedral: [Fe(H₂O)₆]³⁺ Tetrahedral: [CuCl₄]²⁻ Square planar: [Pt(NH₃)₂Cl₂]
74
What are the important transition metal redox systems?
| System | Redox pair | | Iron | Fe³⁺(aq) / Fe²⁺(aq) | | Manganese | MnO₄⁻(aq) / Mn²⁺(aq) | | Chromium | Cr₂O₇²⁻(aq) / Cr³⁺(aq) |
75
What is ligand exchange?
The process where one ligand in a complex ion is replaced by another. Can cause a color change or change in stability of the complex.
76
What is a ligand
A ligand is an atom or atom group that can donate a lone pair of electrons to a transition metal ion to form a complex through the formation of co-ordinate bonds.
77
A complex ion
A complex ion is a molecule or ion where a central metal ion is bonded to one or more ligands
78
Give examples of ligand exchange reactions.
[Cu(H₂O)₆]²⁺ + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O (color changes from blue to deep blue) [Co(H₂O)₆]²⁺ + 4Cl⁻ → [CoCl₄]²⁻ + 6H₂O (pink to blue)
79
How does ligand exchange relate to haemoglobin?
O₂ binds reversibly to the Fe²⁺ center in haemoglobin. CO binds irreversibly, preventing oxygen transport, causing poisoning.
80
Which cations are identified by flame tests?
K⁺: Lilac flame Na⁺: Yellow flame Ca²⁺: Brick red flame Ba²⁺: Pale green flame Cu²⁺: Blue-green flame
81
What principle explains flame test colours?
Atomic emission spectra: Electrons absorb energy, jump to higher levels, then release light when falling back.
82
How do you identify Mg²⁺(aq), Al³⁺(aq), Ca²⁺(aq), Cr³⁺(aq), Mn²⁺(aq), Fe²⁺(aq), Fe³⁺(aq), Cu²⁺(aq), Zn²⁺(aq), Ba²⁺(aq), Pb²⁺(aq), and NH₄⁺(aq)?
Add OH⁻(aq) or NH₃(aq) to form precipitates. Confirmatory tests are needed for some ions.
83
Describe the reactions of common cations with NaOH(aq).
| Cation | Reaction with NaOH(aq) | | Mg²⁺ | White ppt, insoluble in excess | | Al³⁺ | White ppt, soluble in excess | | Ca²⁺ | White ppt, slightly soluble in excess | | Cr³⁺ | Grey-green ppt, soluble in excess | | Mn²⁺ | Buff ppt, oxidizes to brown MnO₂ | | Fe²⁺ | Green ppt, oxidizes to brown | | Fe³⁺ | Brown ppt | | Cu²⁺ | Blue ppt | | Zn²⁺ | White ppt, soluble in excess | | Ba²⁺ | No ppt or very slight white ppt | | Pb²⁺ | White ppt, soluble in excess |
84
How does NH₃(aq) react with common cations?
| Mg²⁺ | White ppt, insoluble | Al³⁺ | White ppt, insoluble | | Ca²⁺ | No ppt | | Cr³⁺ | Grey-green ppt, partially soluble | Mn²⁺ | Light brown ppt | | Fe²⁺ | Green ppt, oxidizes | Fe³⁺ | Brown ppt | Cu²⁺ | Blue ppt, dissolves in excess (deep blue solution) | Zn²⁺ | White ppt, soluble in excess | Ba²⁺ | No ppt | Pb²⁺ | White ppt, insoluble | Cation | Reaction with NH₃(aq)
85
What is the confirmatory test for NH₄⁺(aq)?
Add NaOH(aq), warm gently, test gas with damp red litmus → turns blue (NH₃ gas evolved).
86
What concepts explain cation precipitation reactions?
Equilibrium concepts: Solubility product (Ksp) and ionic product. Basic, amphoteric oxides: Al(OH)₃, Zn(OH)₂ show amphoterism. Complexation: Ligand exchange stabilizes some precipitates in excess OH⁻ or NH₃.
87
Write an ionic equation example for a hydroxide precipitation.
Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) (state symbols must be included!)
88
Which anions are identified in qualitative analysis?
CO₃²⁻ (carbonate) NO₃⁻ (nitrate) SO₄²⁻ (sulfate) SO₃²⁻ (sulfite) Cl⁻ (chloride) Br⁻ (bromide) I⁻ (iodide) CrO₄²⁻ (chromate)
89
How do you test for CO₃²⁻?
Add dilute HCl(aq). Effervescence (CO₂ gas). Confirm with limewater: Ca(OH)₂(aq) turns milky. Equation: CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)
90
How do you test for SO₄²⁻ and SO₃²⁻?
Add Ba²⁺(aq) (e.g., BaCl₂(aq)): SO₄²⁻: White ppt of BaSO₄ (insoluble in dilute acid). SO₃²⁻: White ppt of BaSO₃ (dissolves in dilute acid).
91
How do you test for NO₃⁻ (nitrate)?
Add Al or Zn powder in alkaline solution, warm. NH₃ gas released (test with damp red litmus → blue). OR Copper turnings + conc. H₂SO₄: brown NO₂ gas evolves.
92
# [](http://) How do you test for Cl⁻, Br⁻, and I⁻?
Add dilute HNO₃ then AgNO₃: Cl⁻: White ppt (AgCl). Br⁻: Cream ppt (AgBr). I⁻: Yellow ppt (AgI). Confirm with aqueous NH₃: AgCl dissolves in dilute NH₃. AgBr dissolves only in conc. NH₃. AgI is insoluble.
93
How do you test for CrO₄²⁻ (chromate)?
Add lead(II) nitrate: yellow ppt of PbCrO₄ forms.
94
Write an ionic equation for a halide reaction with AgNO₃.
**Ag⁺(aq) + Cl⁻(aq) → AgCl(s)** AgNO3(aq) + KCl(aq) → AgCl + KNO3(s)
95
indicator pH range summary
| **Indicator** | **pH Range** | **Color Change (Acid → Base)** | **Best for Titration Type** | | **Methyl Orange** | 3.1 – 4.4 | Red → Yellow | **Strong acid – Weak base** | | **Bromophenol Blue** | 3.0 – 4.6 | Yellow → Blue | Strong acid – Weak base | | **Litmus** | \~4.5 – 8.3 | Red → Blue | General acid/base test (not sharp endpoint) | | **Methyl Red** | 4.2 – 6.2 | Red → Yellow | Strong acid – Weak base | | **Bromothymol Blue** | 6.0 – 7.6 | Yellow → Blue | **Strong acid – Strong base** | | **Phenol Red** | 6.4 – 8.4 | Yellow → Red | Strong acid – Strong base | | **Phenolphthalein** | 8.2 – 10.0 | Colourless → Pink | **Weak acid – Strong base**, also strong base |
96
weak vs strong
partially ionise fully dissociate

CAPE Chemistry (5 decks)

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