IGCSE Chemistry Paper 4 - Topic 2: Atoms, Elements and Compounds
Topic 2: Atoms, Elements and Compounds -- Challenge Exam
75 minutes
80
7
75:00
⚡ Cambridge Challenge Level
These questions match real Cambridge IGCSE difficulty. Scoring 50%+ is a solid B, 60%+ is an A, 70%+ is A*. Don't worry if this feels harder -- that's the point!
Instructions
Answer ALL questions in the spaces provided.
Show all working for calculation questions. Marks are awarded for method even if the final answer is incorrect.
Where explanations are required, use correct scientific terminology.
Your answers will be automatically graded when you submit the exam. The model answers will be shown for review.
Question Navigation Score: 0 / 80
Question 1: Elements, Compounds, Mixtures and Atomic Structure
Total: 12 marks
A research team at the Indian Institute of Science in Bangalore analyses four samples labelled W, X, Y and Z. Their findings are summarised in the table below.
Sample
Melting point
Can it be separated by physical methods?
Number of types of atom present
W
Sharp (fixed)
No
1
X
Sharp (fixed)
No
3
Y
Range
Yes
2
Z
Sharp (fixed)
No
2
(a)3 marks
Classify each sample as an element, compound or mixture. In each case, explain how the data in the table supports your answer.
Model Answer - Q1(a)
W is an element -- it has a sharp melting point (pure), cannot be separated by physical methods, and contains only one type of atom [1]
X is a compound -- it has a sharp melting point (pure), cannot be separated by physical methods (chemically bonded), and contains three types of atom [1]
Y is a mixture -- it melts over a range (impure), can be separated by physical methods, and contains two types of atom that are not chemically bonded. Z is a compound -- sharp melting point, cannot be separated physically, contains two types of atom chemically combined [1]
⚠ If you missed marks here: The classic slip is calling Z a mixture just because it has 2 types of atom – number of atom types alone never decides it; Z has a sharp melting point and cannot be separated physically, so it must be a compound. You also lose the mark if you classify correctly but do not quote the table data as evidence – each mark needs classification PLUS the supporting data.
(b)4 marks
A student states: "All pure substances are elements." Evaluate this statement using examples from the table and your own knowledge. Explain the key differences between elements and compounds.
Model Answer - Q1(b)
The statement is incorrect -- compounds such as X and Z are also pure substances because they have sharp melting points [1]
An element contains only one type of atom and cannot be broken down into simpler substances by chemical means [1]
A compound contains two or more types of atom chemically combined in a fixed ratio [1]
Compounds can be broken down into elements by chemical reactions but not by physical methods, whereas elements cannot be broken down at all [1]
⚠ If you missed marks here: Answers that only say "the statement is wrong" without naming X or Z as pure compounds lose the first mark – an evaluate question needs evidence. The other common losses are writing "a compound is made of different atoms" without saying chemically combined in a fixed ratio, and forgetting the chemical-vs-physical breakdown point entirely.
(c)5 marks
Sample W is identified as potassium (K). The most common isotope of potassium has 19 protons and 20 neutrons.
(i) State the proton number and the nucleon number of this isotope. (2 marks)
(ii) Write the full electronic configuration of a potassium atom. (1 mark)
(iii) Explain how the electronic configuration of potassium relates to its position in the Periodic Table. (2 marks)
Model Answer - Q1(c)
(i) Proton number (Z) = 19 [1]
(i) Nucleon number (A) = 19 + 20 = 39 [1]
(ii) Electronic configuration: 2, 8, 8, 1 [1]
(iii) Potassium has 4 electron shells so it is in Period 4 [1]
(iii) Potassium has 1 electron in its outer shell so it is in Group I [1]
⚠ If you missed marks here: If you wrote nucleon number = 20 you gave the neutron count – nucleon number is protons + neutrons = 19 + 20 = 39. For the configuration, 2, 8, 9 is the classic error: the third shell holds only 8 before the fourth starts at IGCSE, so it must be 2, 8, 8, 1. In (iii) you need BOTH links – 4 shells → Period 4 AND 1 outer electron → Group I – stating just one earns just one mark.
Question 2: Atomic Structure, Electronic Configuration and the Periodic Table
Total: 12 marks
The table below shows information about six particles, A to F. Some values are missing.
Particle
Proton number
Nucleon number
Number of electrons
Charge
A
11
23
11
0
B
11
23
10
+1
C
17
35
18
-1
D
12
24
?
+2
E
?
40
18
+2
F
16
32
?
-2
(a)3 marks
Determine the missing values for particles D, E and F. Show your reasoning in each case.
Model Answer - Q2(a)
D: charge = +2 means it has lost 2 electrons, so electrons = 12 - 2 = 10 [1]
E: charge = +2 means protons exceed electrons by 2, so proton number = 18 + 2 = 20 (calcium) [1]
F: charge = -2 means it has gained 2 electrons, so electrons = 16 + 2 = 18 [1]
⚠ If you missed marks here: The usual mistake is flipping the sign logic: a +2 charge means the particle has LOST 2 electrons (fewer electrons than protons), so D = 12 – 2 = 10, not 14. If you got E = 16 you subtracted instead of added – for a +2 ion the protons must EXCEED the 18 electrons by 2, giving 20. Charge = protons – electrons is the safe check every time.
(b)3 marks
(i) Identify the relationship between particles A and B. Explain your answer. (2 marks)
(ii) Which two particles in the table have the same electronic configuration? What term describes particles with the same electronic configuration? (1 mark)
Model Answer - Q2(b)
(i) A is a sodium atom and B is a sodium ion (Na+) -- they have the same proton number and nucleon number [1]
(i) B has lost one electron to form a cation with a +1 charge, while A is electrically neutral [1]
(ii) B, D, E, C and F all have 18 electrons but B and D both have 10 electrons (same as neon). Particles with the same electronic configuration are called isoelectronic [1]
⚠ If you missed marks here: Many students call A and B "isotopes" – wrong, because isotopes differ in NEUTRONS, but A and B have identical proton and nucleon numbers and differ only in electrons: B is the Na+ ion formed when atom A loses one electron. In (ii), saying "B and D are the same" without the term isoelectronic (or without spotting they both have 10 electrons like neon) drops the mark.
(c)4 marks
An atom of element G has the electronic configuration 2, 8, 7.
(i) State the proton number, group and period of element G. Identify the element. (2 marks)
(ii) Element G forms an ion. State the charge on this ion and write its electronic configuration. Explain why G forms an ion with this charge. (2 marks)
Model Answer - Q2(c)
(i) Proton number = 2 + 8 + 7 = 17, Group VII, Period 3, element is chlorine (Cl) [1]
(i) The number of outer electrons equals the group number, the number of occupied shells equals the period number [1]
(ii) Chlorine gains 1 electron to form Cl- with charge -1 and electronic configuration 2, 8, 8 [1]
(ii) It gains 1 electron to achieve a full outer shell, which is the stable electronic configuration of the nearest noble gas (argon) [1]
⚠ If you missed marks here: The proton number is just 2 + 8 + 7 = 17 – students who guess from the Periodic Table without adding the shells often slip. Writing the ion charge as +1 is the big one: chlorine GAINS an electron (7 outer electrons, so gaining 1 is easier than losing 7), making Cl− with configuration 2, 8, 8. "To be stable" alone is not enough for the last mark – you must say full outer shell / noble gas (argon) configuration.
(d)2 marks
State the relative charge and relative mass of a proton, a neutron and an electron. Present your answer as a table.
Model Answer - Q2(d)
Proton: relative charge = +1, relative mass = 1. Neutron: relative charge = 0, relative mass = 1 [1]
Electron: relative charge = -1, relative mass = 1/1836 or negligible or very small [1]
⚠ If you missed marks here: The mark most often lost is the electron's mass – writing 0 exactly, or "1" – the accepted answers are 1/1836, negligible, or very small. Also check you gave RELATIVE values, not real masses in kg, and that the neutron's charge is 0, not "neutral charge of 1".
Question 3: Isotopes
Total: 12 marks
A mass spectrometer at the University of Oxford is used to analyse samples of different elements.
(a)2 marks
Define the term isotopes. Give an example using nuclide notation.
Model Answer - Q3(a)
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons [1]
Example: carbon-12 (6 protons, 6 neutrons) and carbon-14 (6 protons, 8 neutrons), written as ¹²C and ¹&sup4;C, or any other valid example [1]
⚠ If you missed marks here: "Atoms with different masses" or "different nucleon numbers" alone is incomplete – the definition must state SAME number of protons but DIFFERENT number of neutrons (same element is implied by same protons). For the example mark you need proper nuclide notation like ¹²C and ¹&sup4;C, not just "carbon-12 and carbon-14" written in words with no numbers attached to the symbol.
(b)2 marks
Explain why isotopes of the same element have the same chemical properties but different physical properties.
Model Answer - Q3(b)
Same chemical properties because isotopes have the same number of electrons and the same electronic configuration, and chemical reactions depend on electrons [1]
Different physical properties (e.g. density, rate of diffusion) because they have different numbers of neutrons giving different masses [1]
⚠ If you missed marks here: "Same chemical properties because they are the same element" is circular and scores nothing – you must say same number of electrons / same electronic configuration, because chemical reactions depend on electrons. Likewise, for the physical mark, "different neutrons" alone is not enough: you must link it to different MASSES (affecting density, diffusion rate, etc.).
(c)4 marks
Magnesium has three naturally occurring isotopes. The mass spectrometer data is shown below.
Isotope
Mass number
Relative abundance / %
²&sup4;Mg
24
78.99
²&sup5;Mg
25
10.00
²&sup6;Mg
26
11.01
(i) Calculate the relative atomic mass of magnesium. Give your answer to 2 decimal places. Show all working. (3 marks)
(ii) Explain why the relative atomic mass of magnesium is not a whole number. (1 mark)
Model Answer - Q3(c)
(i) A_r = (24 x 78.99 + 25 x 10.00 + 26 x 11.01) / 100 [1]
(i) = (1895.76 + 250.00 + 286.26) / 100 [1]
(i) = 2432.02 / 100 = 24.32 [1]
(ii) The relative atomic mass is a weighted average of the masses of all the isotopes, taking into account their abundances, so it is not a whole number [1]
⚠ If you missed marks here: If you got 25.00 you took a simple average of 24, 25 and 26 – you must WEIGHT each mass by its abundance: (24 × 78.99 + 25 × 10.00 + 26 × 11.01) ÷ 100 = 24.32. Dividing by 3 instead of 100 is the other classic slip. In (ii), "because of isotopes" is not enough – say it is a weighted AVERAGE of isotope masses accounting for their abundances.
(d)4 marks
Silicon has three stable isotopes: ²&sup8;Si, ²&sup9;Si and ³&sup0;Si. The relative atomic mass of silicon is 28.09. The abundance of ²&sup8;Si is 92.2% and the abundance of ³&sup0;Si is 3.1%.
Calculate the percentage abundance of ²&sup9;Si. Show all working clearly.
Check: A_r = (28 x 92.2 + 29 x 4.7 + 30 x 3.1) / 100 [1]
= (2581.6 + 136.3 + 93.0) / 100 [1]
= 2810.9 / 100 = 28.109 which rounds to 28.11 (close to 28.09, slight difference due to rounding in the given data) [1]
⚠ If you missed marks here: Students often panic and set up algebra with an unknown x, missing the simple starting point: all abundances must add to 100%, so ²&sup9;Si = 100 – 92.2 – 3.1 = 4.7%. The verification marks are then lost by not substituting back into the Aᵣ formula – "show all working" means you must demonstrate 4.7% reproduces roughly 28.09, not just state the answer.
Question 4: Ionic Bonding
Total: 12 marks
Sodium oxide (Na&sub2;O) and magnesium chloride (MgCl&sub2;) are both ionic compounds used in different industrial applications. Sodium oxide is used in glass manufacturing in Firozabad, India, while magnesium chloride is used for de-icing roads in the United Kingdom during winter.
(a)2 marks
Define the term ionic bond. State how a cation and an anion are formed.
Model Answer - Q4(a)
An ionic bond is the electrostatic attraction between oppositely charged ions [1]
A cation is formed when a metal atom loses electrons (positive ion). An anion is formed when a non-metal atom gains electrons (negative ion) [1]
⚠ If you missed marks here: Defining the ionic bond as "the transfer of electrons" loses the mark – transfer is how the ions FORM; the bond itself is the electrostatic attraction between oppositely charged ions. The other frequent error is swapping the words: a CATion is POSitive (loses electrons), an anion is negative (gains electrons).
(b)4 marks
Draw dot-and-cross diagrams to show the formation of:
(i) sodium oxide, Na&sub2;O (2 marks)
(ii) magnesium chloride, MgCl&sub2; (2 marks)
Show the outer electrons only. Show the charges on the ions formed.
Model Answer - Q4(b)
(i) Na&sub2;O: Two Na atoms each lose 1 electron to form Na+ ions (2,8). The O atom gains 2 electrons to form O²- ion (2,8). Dot-and-cross shows Na electrons as dots transferred to O shown as crosses (or vice versa) [1]
(i) Correct charges shown: two Na+ and one O²-, all ions have complete outer shells [1]
(ii) MgCl&sub2;: Mg atom loses 2 electrons to form Mg²+ ion (2,8). Each Cl atom gains 1 electron to form Cl- ion (2,8,8). Dot-and-cross shows Mg electrons transferred one to each Cl [1]
(ii) Correct charges shown: one Mg²+ and two Cl-, all ions with complete outer shells [1]
⚠ If you missed marks here: The most common losses: drawing only ONE Na (Na&sub2;O needs TWO Na+ per O²−) or only one Cl (MgCl&sub2; needs TWO Cl− per Mg²+), and forgetting the charges on the ions – a diagram without +, 2+, −, 2− labels drops a mark every time. Also check you drew IONS after transfer (Na+ shown as 2,8 with an empty outer shell, O²− with 8 outer electrons) and used dots for one element, crosses for the other.
(c)3 marks
Describe the structure of an ionic compound such as sodium chloride. Use this to explain why ionic compounds have high melting points.
Model Answer - Q4(c)
Ionic compounds have a giant ionic lattice structure -- a regular three-dimensional arrangement of alternating positive and negative ions [1]
There are strong electrostatic forces of attraction between oppositely charged ions acting in all directions throughout the lattice [1]
A large amount of energy is needed to overcome these many strong electrostatic attractions, so the melting point is high [1]
⚠ If you missed marks here: Vague answers like "the bonds are strong" miss the marks – you need the full chain: giant lattice of alternating + and − ions, STRONG ELECTROSTATIC forces between OPPOSITELY CHARGED ions acting in all directions, and lots of ENERGY needed to overcome them. Never mention "molecules" or "intermolecular forces" for NaCl – ionic compounds have no molecules, and that word can cost you the mark.
(d)3 marks
Explain why sodium chloride does not conduct electricity when solid but does conduct when molten or dissolved in water.
Model Answer - Q4(d)
In solid sodium chloride, the ions are held in fixed positions in the lattice and cannot move, so they cannot carry charge [1]
When molten, the ionic lattice breaks down and the ions are free to move throughout the liquid [1]
When dissolved in water, the ions are dissociated and are free to move through the solution. In both cases, the mobile ions can carry charge and so conduct electricity [1]
⚠ If you missed marks here: The biggest error is saying ELECTRONS move – in molten or dissolved NaCl it is the IONS that move and carry the charge; electrons only conduct in metals and graphite. Also, "the ions vibrate" in the solid is not the point – you must say they are held in FIXED positions in the lattice and CANNOT MOVE, then say they become free to move when molten or dissolved.
Question 5: Covalent Bonding
Total: 10 marks
Carbon dioxide (CO&sub2;), ethene (C&sub2;H&sub4;) and nitrogen (N&sub2;) are all simple covalent molecules. These substances are important in different industrial processes across the world.
(a)2 marks
Define a covalent bond. State the difference between a single covalent bond and a double covalent bond.
Model Answer - Q5(a)
A covalent bond is a shared pair of electrons between two atoms [1]
A single covalent bond involves one shared pair of electrons, while a double covalent bond involves two shared pairs of electrons [1]
⚠ If you missed marks here: "Sharing electrons" is not precise enough – the definition is a shared PAIR of electrons between two atoms. And a double bond is TWO SHARED PAIRS (4 electrons), not "a stronger bond" or "sharing two electrons" – two electrons is only one pair, i.e. a single bond.
(b)4 marks
Draw dot-and-cross diagrams for the following molecules, showing all outer shell electrons:
(i) carbon dioxide, CO&sub2; (2 marks)
(ii) nitrogen, N&sub2; (2 marks)
Model Answer - Q5(b)
(i) CO&sub2;: C shares 2 pairs of electrons with each O atom (double bonds). C has 4 bonding pairs (no lone pairs). Each O has 2 bonding pairs and 2 lone pairs. Total 16 outer electrons shown [1]
(i) Dots used for one atom and crosses for the other, all electrons in the outer shells correctly placed [1]
(ii) N&sub2;: Each N shares 3 pairs of electrons with the other N (triple bond). Each N has 3 bonding pairs and 1 lone pair. Total 10 outer electrons shown [1]
(ii) Dots and crosses correctly distinguished, showing 3 shared pairs between N atoms plus 1 lone pair on each N [1]
⚠ If you missed marks here: The lone pairs are the classic mark-loser: each O in CO&sub2; must show TWO lone pairs as well as its two bonding pairs, and each N in N&sub2; must show ONE lone pair beside the triple bond – diagrams with only the shared electrons drop a mark per molecule. Also check CO&sub2; has DOUBLE bonds (2 shared pairs each side, not single bonds) and N&sub2; a TRIPLE bond (3 shared pairs) – count that every atom ends up with 8 outer electrons.
(c)4 marks
CO&sub2;, C&sub2;H&sub4; and N&sub2; are all gases at room temperature with low boiling points.
(i) Explain, in terms of structure and bonding, why these simple covalent molecules have low boiling points. (3 marks)
(ii) Carbon dioxide does not conduct electricity. Explain why. (1 mark)
Model Answer - Q5(c)
(i) These are simple molecular substances consisting of small discrete molecules [1]
(i) The covalent bonds within each molecule are strong, but the intermolecular forces (forces between the molecules) are weak [1]
(i) Only a small amount of energy is needed to overcome these weak intermolecular forces to separate the molecules, so the boiling points are low [1]
(ii) CO&sub2; does not conduct because it has no free electrons or ions to carry charge -- all electrons are localised in covalent bonds [1]
⚠ If you missed marks here: The number one misconception in all of IGCSE bonding: saying "the covalent bonds are weak" or "the covalent bonds break on boiling". The covalent bonds are STRONG and stay intact – it is the weak INTERMOLECULAR forces BETWEEN molecules that are overcome, needing little energy, hence low boiling points. For (ii), "it is covalent so it doesn't conduct" scores zero – you must say there are no free electrons or ions to carry charge.
Question 6: Giant Covalent Structures and Comparison
Total: 12 marks
Diamond, graphite and silicon dioxide (SiO&sub2;) are all giant covalent structures. A materials engineering firm in Sheffield, UK, uses these substances in cutting tools, lubricants and electronics.
(a)4 marks
Compare the structures of diamond and graphite. For each, state the number of covalent bonds each carbon atom forms and describe the overall arrangement.
Model Answer - Q6(a)
In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement [1]
This forms a rigid three-dimensional giant covalent structure with no free electrons [1]
In graphite, each carbon atom is covalently bonded to three other carbon atoms in flat hexagonal layers [1]
The fourth outer electron on each carbon is delocalised between the layers. The layers are held together by weak intermolecular forces (van der Waals forces) [1]
⚠ If you missed marks here: A "compare" question needs BOTH numbers stated precisely: 4 bonds per carbon (tetrahedral, 3D) in diamond, 3 bonds per carbon (flat hexagonal layers) in graphite – mixing these up or omitting "tetrahedral"/"layers" costs marks. The mark most often forgotten is the fourth: graphite's spare electron is DELOCALISED and the layers are held by WEAK forces – without both details you cap at 3/4.
(b)4 marks
Using your knowledge of structure and bonding, explain the following observations:
(i) Diamond is the hardest naturally occurring substance, while graphite is soft and slippery. (2 marks)
(ii) Graphite conducts electricity but diamond does not. (2 marks)
Model Answer - Q6(b)
(i) Diamond is hard because each carbon is bonded to four others by strong covalent bonds in all directions throughout the structure, making it very rigid [1]
(i) Graphite is soft because the layers are held together only by weak intermolecular forces, so the layers can slide over each other easily [1]
(ii) In graphite, each carbon uses only 3 of its 4 outer electrons for bonding. The fourth electron is delocalised and free to move along the layers, carrying charge [1]
(ii) In diamond, all four outer electrons on each carbon are used in covalent bonds, so there are no free electrons to carry charge [1]
⚠ If you missed marks here: For softness, "graphite has weak bonds" loses the mark – the covalent bonds WITHIN each layer are strong; it is the weak forces BETWEEN layers that let the LAYERS SLIDE over each other. For conductivity, the answer must be electron-based: graphite's fourth electron per carbon is delocalised and moves along the layers carrying charge, while in diamond all 4 outer electrons are locked in bonds – saying "graphite has ions" or "diamond is too hard to conduct" scores zero.
(c)4 marks
Silicon dioxide (SiO&sub2;) is used in the manufacture of optical fibres and electronics. It has a giant covalent structure similar to diamond.
(i) Describe the structure of silicon dioxide. (2 marks)
(ii) Silicon dioxide has a very high melting point (1710 °C) and does not conduct electricity. Explain both of these properties in terms of its structure and bonding. (2 marks)
Model Answer - Q6(c)
(i) SiO&sub2; has a giant covalent structure where each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms [1]
(i) This forms a continuous three-dimensional network of strong covalent bonds (similar to diamond but with alternating Si and O atoms) [1]
(ii) High melting point because a very large amount of energy is needed to break the many strong covalent bonds throughout the giant structure [1]
(ii) Does not conduct electricity because all outer electrons are used in covalent bonds, so there are no free electrons or ions to carry charge [1]
⚠ If you missed marks here: The structure mark needs the exact ratio: each Si bonded to FOUR O atoms and each O bonded to TWO Si atoms – "silicon bonded to oxygen" without the numbers is too vague. For the 1710 °C melting point, the energy must go into breaking many STRONG COVALENT BONDS throughout the giant structure – if you wrote "intermolecular forces" you treated SiO&sub2; as a simple molecule, which is the classic confusion.
Question 7: Metallic Bonding
Total: 10 marks
Copper is one of the most widely used metals in the world. The Khetri copper mines in Rajasthan, India, produce copper that is used for electrical wiring, water pipes and cooking vessels. Understanding metallic bonding helps explain why copper is suitable for these applications.
(a)3 marks
Describe metallic bonding. Include the terms delocalised electrons, positive ions and electrostatic attraction in your answer.
Model Answer - Q7(a)
In a metal, the atoms lose their outer shell electrons to form positive ions (cations) arranged in a regular lattice [1]
The outer electrons become delocalised -- they are free to move throughout the entire metal structure, forming a "sea" of electrons [1]
Metallic bonding is the strong electrostatic attraction between the positive metal ions and the delocalised electrons surrounding them [1]
⚠ If you missed marks here: Writing "positive ATOMS" instead of positive IONS is the top error – the atoms have lost their outer electrons, so a lattice of positive ions remains. The question told you the three required terms, so check all three appear: sea of DELOCALISED ELECTRONS, POSITIVE IONS in a lattice, and ELECTROSTATIC ATTRACTION between them – each missing term is usually a missing mark.
(b)4 marks
Using the model of metallic bonding, explain the following properties of metals:
(i) Metals are good conductors of electricity. (2 marks)
(ii) Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires). (2 marks)
Model Answer - Q7(b)
(i) Metals contain delocalised electrons that are free to move through the structure [1]
(i) When a potential difference is applied, these electrons can drift in one direction, carrying charge through the metal as an electric current [1]
(ii) When a force is applied, the layers of positive ions can slide over each other into new positions [1]
(ii) The metallic bonding is maintained because the delocalised electrons can adjust their positions around the ions in the new arrangement. The bonding is non-directional so the structure does not shatter [1]
⚠ If you missed marks here: In metals the charge carriers are delocalised ELECTRONS, not ions – and "electrons are free" alone only earns the first mark; the second needs them to MOVE/drift carrying charge when a potential difference is applied. For malleability, "layers of ions slide" gets one mark, but most students miss the second: the bonding is MAINTAINED because the delocalised electrons still surround the ions in the new positions, so the metal bends instead of shattering.
(c)3 marks
The table below summarises the bonding and properties of four different types of substance. Complete the table by filling in the missing information for substance types 3 and 4.
Type 1: Ionic compound
Type 2: Simple covalent
Type 3: Giant covalent
Type 4: Metal
Example
NaCl
CO&sub2;
?
?
Melting point
High
Low
?
?
Conducts electricity?
Only when molten/dissolved
No
?
?
Model Answer - Q7(c)
Type 3 (Giant covalent): Example = diamond or SiO&sub2;, Melting point = Very high, Conducts = No (except graphite which conducts due to delocalised electrons) [1]
Type 4 (Metal): Example = copper or iron or any named metal, Melting point = Generally high (varies) [1]
Type 4 (Metal): Conducts electricity = Yes, in both solid and liquid states (due to delocalised electrons) [1]
⚠ If you missed marks here: For giant covalent, "No" without the graphite exception can drop the mark – diamond and SiO&sub2; do not conduct, but graphite does (delocalised electrons), so mention it. For metals, the key contrast with ionic compounds is that they conduct when SOLID as well as molten – writing "only when molten" copies the ionic column and loses the mark.
Exam Score Summary
0
out of 80
0%
-
A* : 56+
A : 48-55
B : 40-47
C : 32-39
D : 24-31
E : 16-23
U : <16
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