inorganic I Flashcards
(111 cards)
1
Q
A
2
Q
What is periodicity?
A
The repeating trends in physical and chemical properties of elements across different periods in the periodic table.
3
Q
What are periods in the periodic table?
A
Horizontal rows in the periodic table.
4
Q
What are groups in the periodic table?
A
Vertical columns containing elements with similar chemical properties.
5
Q
Why do atomic radius and other properties show periodic trends?
A
Due to the increasing nuclear charge and changes in electron shielding across periods.
6
Q
How does atomic radius change across a period?
A
Atomic radius decreases across a period due to increased nuclear charge pulling electrons closer.
7
Q
How does atomic radius change down a group?
A
Atomic radius increases down a group due to added electron shells.
8
Q
Why does the first ionisation energy generally increase across a period?
A
Because of increasing nuclear charge and decreasing atomic radius, making it harder to remove an electron.
9
Q
Why does the first ionisation energy generally decrease down a group?
A
Because increased electron shielding and larger atomic radius make it easier to remove an electron.
10
Q
What causes exceptions (drops) in ionisation energy across a period?
A
Electron repulsion in paired electrons (e.g. between groups 2 and 3, and groups 5 and 6).
11
Q
How does metallic reactivity change across a period?
A
Metallic character decreases across a period.
12
Q
How does metallic reactivity change down a group?
A
Metallic character increases down a group.
13
Q
What are the typical properties of metals?
A
Conductive, malleable, ductile, tend to lose electrons to form positive ions.
14
Q
What are the typical properties of non-metals?
A
Poor conductors, brittle if solid, tend to gain electrons to form negative ions.
15
Q
Describe the trend in melting points across Period 3 elements.
A
Increase from Na to Si (metallic bonding gets stronger), then sharp decrease from P to Ar (simple molecular structures with weak intermolecular forces).
16
Q
Why does Silicon have a high melting point?
A
Because of its giant covalent structure with strong covalent bonds.
17
Q
How does reactivity change across Period 3 for metals?
A
Reactivity decreases across the metals (Na to Mg to Al) due to increasing ionisation energy.
18
Q
How does reactivity change across Period 3 for non-metals?
A
Reactivity increases across non-metals (P to Cl) due to increasing electronegativity.
19
Q
How do the oxides of Period 3 elements change across the period?
A
Oxides change from basic (Na₂O) to amphoteric (Al₂O₃) to acidic (P₄O₁₀, SO₂, SO₃).
20
Q
What is the acid-base character of Na₂O?
A
Basic oxide.
21
Q
What is the acid-base character of Al₂O₃?
A
Amphoteric oxide.
22
Q
What is the acid-base character of SO₃?
A
Acidic oxide.
23
Q
Write the reaction of sodium oxide with water.
A
Na₂O + H₂O → 2NaOH
24
Q
Write the reaction of sulfur dioxide with water.
A
SO₂ + H₂O → H₂SO₃ (sulfurous acid)
25
Write the reaction of sulfur trioxide with water.
SO₃ + H₂O → H₂SO₄ (sulfuric acid)
26
What type of bonding does sodium oxide have?
Ionic bonding.
27
What type of bonding does silicon dioxide have?
Giant covalent bonding.
28
What type of bonding do phosphorus pentoxide and sulfur oxides have?
Simple molecular covalent bonding.
29
How does the electrical conductivity of silicon dioxide compare to sodium oxide?
Silicon dioxide does not conduct electricity (no free electrons), sodium oxide conducts when molten (ions can move).
30
Equation for sodium oxide reacting with water:
Na₂O + H₂O → 2NaOH
31
Equation for magnesium oxide reacting with water:
MgO + H₂O → Mg(OH)₂
32
Equation for phosphorus pentoxide reacting with water:
P₄O₁₀ + 6H₂O → 4H₃PO₄
33
Equation for sulfur dioxide reacting with water:
SO₂ + H₂O → H₂SO₃
34
Equation for sulfur trioxide reacting with water:
SO₃ + H₂O → H₂SO₄
35
Across Period 3, how do melting points vary?
Increase from Na to Si due to metallic to giant covalent bonding, then sharply drop for P to Ar due to simple molecular structures.
36
Across Period 3, how does metallic character change?
Decreases from left to right.
37
Across Period 3, how does reactivity of metals change?
Decreases.
38
Across Period 3, how does acidity of oxides change?
Oxides become more acidic.
39
What causes periodic trends across a period?
Increasing nuclear charge and similar electron shielding.
40
How does atomic radius change across Period 3?
It decreases due to increasing nuclear charge attracting electrons closer.
41
How does atomic radius change down Group 2?
It increases due to added electron shells.
42
Give reasons for the ionisation energy drop between groups 2 and 3 in Period 3.
Group 3 has electrons in a new p-orbital which is higher energy and more shielded, making it easier to remove.
43
Describe the trend in melting points across Period 3.
Increases from Na to Si due to stronger metallic and covalent bonds, then sharply drops from P to Ar due to weak intermolecular forces.
44
What is the acid-base character of oxides across Period 3?
Oxides change from basic (Na₂O) to amphoteric (Al₂O₃) to acidic (SiO₂, P₄O₁₀, SO₂, SO₃).
45
Explain why atomic radius decreases across Period 3. (4 marks)
Nuclear charge increases as protons increase; electrons added to the same shell so shielding remains similar; increased attraction pulls electrons closer, reducing atomic radius.
46
Explain why magnesium has a higher melting point than sodium. (6 marks)
Magnesium ions have a +2 charge, sodium ions +1; magnesium has more delocalised electrons; stronger metallic bonding in magnesium requires more energy to break bonds, so melting point is higher.
47
Describe and explain the trend in first ionisation energy across Period 3. (6 marks)
First ionisation energy generally increases due to increasing nuclear charge and similar shielding, making it harder to remove electrons; drops between Mg and Al because electron enters higher-energy p-orbital; drops between P and S due to electron pairing repulsion.
48
Write balanced equations for the reaction of phosphorus pentoxide with water. (2 marks)
P₄O₁₀ + 6H₂O → 4H₃PO₄
49
Explain why aluminium oxide is amphoteric. (4 marks)
Aluminium oxide reacts with both acids and bases; it has both ionic and covalent character; reacts with acids to form salts; reacts with bases to form aluminate ions.
50
Explain why silicon dioxide has a high melting point. (3 marks)
It has a giant covalent structure with strong covalent bonds throughout; lots of energy needed to break bonds; resulting in high melting point.
51
Compare the structure and bonding of sodium oxide and phosphorus pentoxide. (6 marks)
Sodium oxide is ionic with Na⁺ and O²⁻ ions; high melting point due to strong ionic bonds; phosphorus pentoxide is molecular with covalent bonds; low melting point due to weak intermolecular forces.
52
Predict the products and write the balanced equation for the reaction of sulfur trioxide with water. (2 marks)
SO₃ + H₂O → H₂SO₄
53
Explain the trend in reactivity of Period 3 metals. (4 marks)
Reactivity decreases across the period as ionisation energy increases; harder to lose electrons; sodium is most reactive; aluminium least reactive among metals in Period 3.
54
Suggest why argon has a very low boiling point. (3 marks)
Argon is a noble gas with monatomic molecules; only weak London dispersion forces between atoms; little energy needed to overcome these forces.
55
Explain why sulfur dioxide is acidic. (3 marks)
It forms sulfurous acid in water; contains polar covalent bonds; releases H⁺ ions making solution acidic.
56
What elements are in Group 2 of the periodic table?
Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra).
57
What is the general electronic configuration of Group 2 elements?
ns² (two electrons in the outer s-subshell).
58
What is the common oxidation state of Group 2 elements in compounds?
2
59
How do atomic radius and ionic radius change down Group 2?
Both increase due to added electron shells.
60
How does first ionisation energy change down Group 2?
Decreases due to increased atomic radius and electron shielding.
61
Describe the trend in melting points down Group 2.
Generally decreases down the group due to weaker metallic bonding (increasing atomic radius means weaker attraction between ions and delocalised electrons).
62
How does reactivity change down Group 2?
Reactivity increases down the group because ionisation energy decreases, making it easier to lose electrons.
63
What type of reactions do Group 2 metals commonly undergo?
Reactions with water, oxygen, and acids to form metal hydroxides, oxides, and salts.
64
Write the general equation for a Group 2 metal reacting with water.
M + 2H₂O → M(OH)₂ + H₂
65
What is observed when magnesium reacts with cold water?
Reacts very slowly with cold water but reacts more rapidly with steam to form magnesium oxide and hydrogen.
66
What is observed when calcium reacts with cold water?
Reacts steadily to form calcium hydroxide (sparingly soluble) and hydrogen gas; solution may turn cloudy due to Ca(OH)₂ precipitate.
67
How does the solubility of Group 2 hydroxides change down the group?
Increases down the group (Mg(OH)₂ is sparingly soluble, Ba(OH)₂ is quite soluble).
68
How does the solubility of Group 2 sulfates change down the group?
Decreases down the group (MgSO₄ is soluble, BaSO₄ is insoluble).
69
Write the equation for the dissolution of calcium hydroxide in water.
Ca(OH)₂ (s) ⇌ Ca²⁺ (aq) + 2OH⁻ (aq)
70
What is calcium hydroxide used for in agriculture?
To neutralise acidic soil (lime treatment).
71
How is barium sulfate used medically?
As a radiopaque agent in barium meals to highlight the gut in X-rays (it is insoluble and non-toxic).
72
Reaction of Group 2 metal with water:
M + 2H₂O → M(OH)₂ + H₂ (only for liquid water = metal hydroxide. If it is steam = metal oxide)
73
Reaction with oxygen:
2M + O₂ → 2MO
74
Formation of sulfates:
MO + H₂SO₄ → MSO₄ + H₂O
75
Dissolution of hydroxides:
M(OH)₂ (s) ⇌ M²⁺ (aq) + 2OH⁻ (aq)
76
Describe the trend in reactivity of Group 2 metals and explain why it occurs. (4 marks)
Reactivity increases down the group as the atomic radius increases; this causes outer electrons to be further from the nucleus and more shielded; thus, less energy is needed to remove electrons.
77
Write an equation for the reaction of magnesium with steam. (2 marks)
Mg + H₂O (g) → MgO + H₂
78
Explain why calcium hydroxide is sparingly soluble in water, but barium hydroxide is more soluble. (4 marks)
Solubility increases down Group 2 because lattice enthalpy decreases faster than hydration enthalpy, making dissolution more favorable for heavier metals.
79
Explain why the solubility of Group 2 sulfates decreases down the group. (4 marks)
Sulfate ions are large and highly charged; lattice enthalpy increases down the group faster than hydration enthalpy; so heavier sulfates are less soluble.
80
Write the balanced equation for the reaction of barium with oxygen. (2 marks)
2Ba + O₂ → 2BaO
81
Explain how calcium hydroxide is used in agriculture. (3 marks)
It neutralises acidic soils by reacting with H⁺ ions, increasing soil pH and improving crop growth.
82
Explain why magnesium reacts differently with cold water compared to calcium. (3 marks)
Magnesium reacts very slowly with cold water due to a protective oxide layer; calcium reacts more readily due to its lower ionisation energy.
83
Predict and explain the trend in melting points of Group 2 metals. (3 marks)
Melting points decrease down the group because the metal ions get larger, reducing the strength of metallic bonding.
84
What is the oxidation state of halogens in their elemental form?
0
85
What is the trend in atomic radius down Group 7?
Increases due to additional electron shells.
86
What is the trend in electronegativity down Group 7?
Decreases — atomic radius increases and shielding reduces nuclear attraction for bonding electrons.
87
What is the trend in boiling points down Group 7?
Increases — due to stronger London dispersion forces with more electrons and larger molecules.
88
What is the trend in reactivity of Group 7 elements?
Decreases down the group — atoms gain electrons less easily due to increased shielding and atomic radius.
89
What is a displacement reaction in Group 7?
A more reactive halogen displaces a less reactive halide ion from solution.
90
What happens when chlorine is added to a solution of potassium iodide?
Chlorine displaces iodine: Cl₂ + 2KI → 2KCl + I₂
91
How can displacement reactions be used to determine halogen reactivity?
Observing colour changes in solutions (e.g., brown for I₂, orange for Br₂) indicates displacement.
92
What is the trend in oxidising ability of halogens?
Decreases down the group — F₂ is the strongest oxidising agent, I₂ is the weakest.
93
In redox reactions, what role do halogens play?
Halogens act as oxidising agents (they gain electrons and are reduced).
94
What is observed when NaCl reacts with H₂SO₄?
Steamy HCl fumes — not a redox reaction: NaCl + H₂SO₄ → NaHSO₄ + HCl
95
What happens when NaBr reacts with H₂SO₄?
Redox occurs — Br⁻ is oxidised, sulfur is reduced: NaBr + H₂SO₄ → NaHSO₄ + HBr and 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O
96
What happens when NaI reacts with H₂SO₄?
Multiple redox steps — I⁻ is a stronger reducing agent: 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O and 6HI + H₂SO₄ → 3I₂ + H₂S + 2H₂O
97
What is disproportionation?
A reaction where the same element is both oxidised and reduced.
98
Give an example of a disproportionation reaction involving chlorine in water.
Cl₂ + H₂O → HCl + HClO (Cl is reduced to HCl and oxidised to HClO)
99
Write the equation for chlorine reacting with cold dilute NaOH.
Cl₂ + 2NaOH → NaCl + NaClO + H₂O (used to make bleach)
100
What is the oxidation state of Cl in Cl₂, NaCl, and NaClO in the above reaction?
0 in Cl₂, –1 in NaCl, +1 in NaClO
101
What reagent is used to test for halide ions?
Acidified silver nitrate (AgNO₃ + HNO₃).
102
What is observed when AgNO₃ is added to Cl⁻, Br⁻ and I⁻?
Cl⁻: White ppt (AgCl); Br⁻: Cream ppt (AgBr); I⁻: Yellow ppt (AgI)
103
What happens when dilute and concentrated ammonia are added to silver halide precipitates?
AgCl: Dissolves in dilute NH₃; AgBr: Dissolves in concentrated NH₃; AgI: Insoluble in NH₃
104
Why is chlorine added to drinking water?
To kill bacteria — forms HOCl, a weak bleach.
105
Give one risk of using chlorine in water treatment.
Chlorinated hydrocarbons (carcinogens) can form when chlorine reacts with organic compounds.
106
Explain the trend in boiling points down Group 7. (3 marks)
Boiling point increases due to more electrons → stronger van der Waals forces → more energy needed to overcome intermolecular forces.
107
Explain the trend in oxidising ability of halogens down the group. (3 marks)
Decreases because atomic radius and shielding increase, so halogens attract electrons less readily.
108
Describe what you would see when chlorine is added to a solution of potassium iodide. (3 marks)
The solution turns brown due to formation of iodine; chlorine displaces iodide.
109
Write the ionic equation for the reaction between chlorine and iodide ions. (1 mark)
Cl₂ + 2I⁻ → 2Cl⁻ + I₂
110
Describe how you would test for a chloride ion in solution. (3 marks)
Add nitric acid, then silver nitrate → white ppt (AgCl); dissolves in dilute NH₃.
111
Chlorine reacts with water in a reversible reaction. Write the equation and explain the benefit of this reaction in water treatment. (4 marks)
Cl₂ + H₂O ⇌ HCl + HOCl; HOCl kills bacteria and acts as a weak bleach without harmful buildup.
