Topic 2
30 questions: 15 multiple choice, 15 structured.
Extended candidates are examined on Core and Supplement content, so keep All.
Which statement correctly describes a compound?
Answer: B
The key test for a compound is two things at once: (1) it contains more than one element, and (2) those elements are chemically joined in a fixed ratio, so a sample of the compound cannot be separated back into its elements by physical methods such as filtering or evaporating - only by a chemical reaction. Checking each option against both tests picks out B.
In which region of an atom are protons and neutrons found?
Answer: A
Picture the atom as a very small, dense nucleus (containing all the protons and neutrons) with electrons moving around it in shells at much greater distances. Almost all of the atom's mass is concentrated in this tiny central nucleus, while the electron shells take up almost all of the atom's volume but almost none of its mass.
Which row correctly gives the relative charge and the relative mass of a neutron?
| Option | Relative charge | Relative mass |
|---|---|---|
| +1 | 1 | |
| -1 | 1 | |
| 0 | very small (about 1/1840) | |
| 0 | 1 |
Answer: D
There are only three particles to remember here, and two facts about each: proton (+1 charge, mass 1), neutron (0 charge, mass 1) and electron (0 charge, mass about 1/1840, i.e. negligible). The neutron is the only particle with both "0" for charge and "1" for mass, which is option D.
An atom of phosphorus has a mass number of 31 and contains 16 neutrons. What is its proton (atomic) number?
Answer: C
Mass number = number of protons + number of neutrons. Rearranging: number of protons = mass number - number of neutrons = 31 - 16 = 15. It helps to keep the two definitions straight: proton number counts only protons; mass number counts protons and neutrons together.
An atom of aluminium has a proton (atomic) number of 13 and a mass (nucleon) number of 27.
State the number of neutrons and the number of electrons in this atom.
Give the electronic configuration of this atom.
State the number of occupied electron shells in this atom, and identify the period of the Periodic Table in which aluminium is found.
Aluminium has proton number 13, so a neutral atom has 13 protons and 13 electrons. The mass number (27) is protons + neutrons, so neutrons = 27 - 13 = 14. To work out the electronic configuration, fill the shells from the nucleus outwards: the first shell holds up to 2 electrons, the second up to 8, and any remaining electrons go into the third shell. With 13 electrons: 2 in the first shell, 8 in the second shell, leaving 13 - 2 - 8 = 3 in the third shell, giving 2,8,3. The number of occupied shells (3) tells you the period number directly - aluminium is in Period 3.
Describe the structure of an atom, including the location of protons, neutrons and electrons, and explain why an atom has no overall electrical charge.
Complete the table to give the relative charge and the relative mass of each particle.
| Particle | Relative charge | Relative mass |
|---|---|---|
| proton | ||
| neutron | ||
| electron |
An atom is built from three particles. The protons and neutrons are packed together in a tiny central nucleus, while the electrons occupy shells around the outside. Protons carry a charge of +1 and electrons carry a charge of -1; because a neutral atom always has equal numbers of protons and electrons, the total positive charge exactly cancels the total negative charge, leaving no overall charge. For the table: protons and neutrons both have a relative mass of 1 (and neutrons have no charge, matching their name), while electrons have almost no mass (about 1/1840 of a proton's mass) and carry a charge of -1.
A sulfur atom (proton number 16) gains two electrons to form a sulfide ion, S2−.
Deduce the number of protons and the number of electrons in the S2− ion, giving a reason for your answer.
Give the electronic configuration of the S2− ion.
A calcium atom (proton number 20) loses two electrons to form a Ca2+ ion. Give the electronic configuration of the Ca2+ ion, and state how many occupied electron shells it has.
Gaining or losing electrons never changes the number of protons in the nucleus - only the number of electrons changes. Sulfur has proton number 16, so the S2− ion still has 16 protons, but it has gained 2 extra electrons, giving 16 + 2 = 18 electrons. Filling the shells with 18 electrons: 2, then 8, then 8, giving 2,8,8. Calcium (proton number 20) loses 2 electrons to form Ca2+, leaving 20 - 2 = 18 electrons - the same number as the sulfide ion above - so the configuration is also 2,8,8, occupying 3 shells. Notice both ions end up with the same electron arrangement as the noble gas argon.
Which statement correctly defines isotopes?
Answer: A
The word "isotope" only ever describes atoms of the same element - same proton number - that differ in their number of neutrons (and therefore in mass number). Any option that mentions "different elements" or "different numbers of protons" cannot be describing isotopes, which immediately rules out B and C.
Chlorine has two common isotopes, 35Cl and 37Cl. Both isotopes have identical chemical properties. Which statement explains why?
Answer: C
Chemical reactions happen through the transfer or sharing of electrons, particularly the outer-shell electrons. Because 35Cl and 37Cl both have 17 protons, they both have 17 electrons arranged in exactly the same way (2,8,7). Since their electron arrangements are identical, their chemical behaviour is identical too - only their masses (and physical properties linked to mass) differ.
Potassium has two isotopes, which can be written as 3919K and 4119K.
For the isotope 3919K, state the number of protons, the number of neutrons and the number of electrons.
Explain, in terms of subatomic particles, why 3919K and 4119K are both isotopes of potassium.
The subscript number (19) in 3919K is the proton number and is the same for both isotopes - this is what makes both of them potassium. The superscript number (39 or 41) is the mass number (protons + neutrons). For 3919K: protons = 19, so neutrons = 39 - 19 = 20, and in a neutral atom electrons = protons = 19. Because both isotopes share the same proton number but differ in mass number, they must have different numbers of neutrons - that is exactly the definition of isotopes.
A sample of chlorine gas contains two isotopes: 75% of the atoms are 35Cl (relative isotopic mass 35) and 25% of the atoms are 37Cl (relative isotopic mass 37).
Calculate the relative atomic mass of this sample of chlorine. Give your answer to 3 significant figures.
The relative atomic mass of chlorine given in the Periodic Table is 35.5. State, with a reason, whether your answer to (a) is consistent with this value.
Relative atomic mass is a weighted average of the isotopic masses, weighted by how abundant each isotope is. Multiply each isotope's mass by its percentage abundance, add the results together, then divide by 100 (because the abundances are percentages): (35 x 75 + 37 x 25) / 100 = (2625 + 925) / 100 = 3550 / 100 = 35.5. This matches the value of 35.5 given for chlorine in the Periodic Table, so the calculation is consistent with it. This also shows why chlorine's relative atomic mass is not a whole number: it is an average of two isotopes with different masses.
A magnesium atom loses two electrons when it reacts. What type of ion is formed, and what is its charge?
Answer: B
Every electron carries a charge of -1. If a neutral atom loses 2 electrons, it loses 2 units of negative charge, so what remains is 2 units more positive than before: a charge of 2+. A positively charged ion is always called a cation (a useful memory link: "cat-ion" - a cation is "paws-itive"). So losing electrons always gives a cation, and here specifically a 2+ cation.
Which statement correctly defines an ionic bond?
Answer: D
An ionic bond only exists between ions - particles that already carry a full positive or negative charge - and it is the electrostatic (opposite-charges-attract) force that holds them together. This is different from a covalent bond, which involves sharing electrons rather than transferring them to form separate charged ions.
Which pair of properties is typical of an ionic compound such as sodium chloride?
Answer: A
Two properties always go together for ionic compounds: a high melting/boiling point (because breaking apart the strong lattice of ions takes a lot of energy) and conduction of electricity only when the ions can move freely - that is, when molten or dissolved in water, but never when solid, since the ions are then locked in fixed positions in the lattice.
Which statement best describes the structure of solid sodium chloride?
Answer: B
Ionic compounds do not exist as individual molecules; instead, huge numbers of positive and negative ions pack together in a regular, repeating three-dimensional pattern (a giant lattice), with each ion attracted electrostatically to the oppositely charged ions around it in every direction. This is what makes ionic lattices so difficult to break apart, giving ionic compounds their high melting points.
Potassium and fluorine react together to form the ionic compound potassium fluoride, KF. A potassium atom has the electronic configuration 2,8,8,1 and a fluorine atom has the electronic configuration 2,7.
Describe how a potassium atom and a fluorine atom form ions when they react together.
State the electronic configuration of the K+ ion and of the F− ion.
Explain why potassium fluoride is described as having an ionic bond.
Potassium (2,8,8,1) has just one electron in its outer shell, so it is easiest for it to lose that one electron and reach the stable configuration 2,8,8 (the same as argon), becoming a K+ ion. Fluorine (2,7) has seven outer electrons, so it is easiest for it to gain one more electron to complete its outer shell, reaching 2,8 (the same as neon), becoming an F− ion. The electron lost by potassium is the same electron gained by fluorine. The resulting K+ and F− ions carry opposite charges, so they attract each other strongly - this electrostatic attraction between oppositely charged ions is the ionic bond.
State two physical properties of solid sodium chloride that are typical of ionic compounds.
State whether solid sodium chloride conducts electricity, and state whether molten sodium chloride conducts electricity.
Ionic compounds like sodium chloride have high melting and boiling points because a great deal of energy is needed to overcome the strong electrostatic forces holding the giant lattice of ions together. For conductivity, what matters is whether the ions can move: in the solid, the ions are held tightly in fixed positions in the lattice and cannot move to carry charge, so solid sodium chloride does not conduct electricity. Once melted, the ions are free to move throughout the liquid, so molten sodium chloride does conduct electricity.
Magnesium (electronic configuration 2,8,2) reacts with oxygen (electronic configuration 2,6) to form magnesium oxide, MgO.
Describe how a magnesium atom and an oxygen atom form ions in this reaction.
Give the electronic configuration of the Mg2+ ion and of the O2− ion.
Explain, in terms of electron transfer, why the ratio of magnesium ions to oxide ions in magnesium oxide is 1:1.
Magnesium (2,8,2) has two outer electrons, which it loses to reach the stable configuration 2,8 (like neon), forming Mg2+. Oxygen (2,6) has six outer electrons, needing two more to complete its outer shell, so it gains two electrons to reach 2,8, forming O2−. Because electrons cannot simply disappear, the number transferred must balance exactly: the 2 electrons lost by one magnesium atom are gained by one oxygen atom. This means the ions produced, Mg2+ and O2−, have charges that are equal and opposite, so they combine in a simple 1:1 ratio to give overall zero charge in MgO.
Calcium oxide, CaO, is an ionic compound.
Describe the structure of calcium oxide as a giant ionic lattice.
Explain, in terms of structure and bonding, why calcium oxide has a high melting point.
Explain, in terms of structure and bonding, why calcium oxide conducts electricity when molten but not when solid.
Calcium oxide is a giant ionic lattice: enormous numbers of Ca2+ and O2− ions arranged in a regular, repeating pattern, each attracted electrostatically to its oppositely charged neighbours in every direction. Melting this lattice means overcoming all of these strong attractive forces at once, which needs a large amount of energy - this is why calcium oxide has such a high melting point. Electrical conductivity depends on charged particles being able to move: in the solid, the ions are locked into fixed positions and cannot carry charge from place to place, so it does not conduct; once molten, the lattice breaks down and the ions become free to move, so they can now carry charge and the liquid conducts electricity.
What is a covalent bond?
Answer: C
The defining feature of covalent bonding, unlike ionic bonding, is that electrons are shared rather than transferred completely from one atom to another. Two atoms each contribute one electron to a shared pair, and both atoms are attracted to this shared pair, holding them together.
Which pair of properties is typical of a simple molecular compound such as water?
Answer: D
Simple molecular substances consist of individual molecules held together by weak intermolecular forces, so relatively little energy is needed to separate the molecules from each other, giving low melting and boiling points. Because the molecules themselves are electrically neutral and there are no ions or delocalised electrons free to move and carry charge, these substances are poor conductors of electricity.
Methane, CH4, is a simple molecular compound.
State the number of covalent bonds formed by the carbon atom in a molecule of methane, and state the number of electrons shared in each of these bonds.
Describe, in terms of electron arrangement, how the covalent bonds in methane give both the carbon atom and each hydrogen atom a full outer shell of electrons.
Carbon has 4 electrons in its outer shell and needs 4 more to complete it (a full outer shell of 8); each hydrogen atom has 1 electron and needs 1 more to complete its outer shell of 2 (hydrogen's shell only holds 2 electrons). Methane solves both needs at once: the carbon atom forms one covalent bond with each of four hydrogen atoms, and each bond is a shared pair of 2 electrons. Counting the carbon atom's outer shell: its own 4 electrons plus 1 shared electron from each of the 4 hydrogen atoms gives 8 electrons in total - a full outer shell. Counting each hydrogen atom: its own 1 electron plus 1 shared electron from carbon gives 2 electrons - also a full outer shell.
In a molecule of carbon dioxide, CO2, the carbon atom forms a double covalent bond with each of the two oxygen atoms.
Describe, in terms of shared electrons, the covalent bonding between the carbon atom and each oxygen atom in a molecule of carbon dioxide.
In a molecule of nitrogen, N2, the two nitrogen atoms share electrons so that each nitrogen atom has a full outer shell. State the number of electrons shared between the two nitrogen atoms, and name the type of covalent bond formed.
Carbon has 4 outer electrons and needs 4 more; each oxygen atom has 6 outer electrons and needs 2 more. In CO2, the carbon atom forms a double bond with each oxygen atom: it shares 2 of its electrons with each oxygen atom (using all 4 of its outer electrons in total), and each oxygen atom shares 2 of its own electrons back with carbon. Each double bond is therefore 2 shared pairs (4 electrons). To count outer electrons, each atom counts every electron in every shared pair around it. Carbon has 2 shared pairs with each oxygen atom, so 4 shared pairs (8 electrons) in total - a full outer shell. Each oxygen atom has 2 shared pairs with carbon (4 electrons) plus its own 4 unshared electrons, giving 8 - also a full outer shell. So all three atoms achieve a noble gas electronic configuration. In N2, each nitrogen atom has 5 outer electrons and needs 3 more, so the two atoms share 3 pairs of electrons (6 electrons in total) between them - a triple bond - giving each nitrogen atom a full outer shell of 8.
Iodine, I2, is a simple molecular substance.
Explain, in terms of structure and bonding, why simple molecular substances such as iodine have low melting points and low boiling points.
Explain why simple molecular substances are poor conductors of electricity.
State one way in which the covalent bond within a molecule of iodine differs from the forces between neighbouring molecules of iodine.
It is important not to confuse the strong covalent bonds holding the two atoms together within an iodine molecule with the much weaker forces of attraction between separate iodine molecules. When iodine melts or boils, the molecules themselves do not break apart - the covalent bonds stay intact. Instead, only the weak forces between neighbouring molecules are overcome, which takes relatively little energy, giving iodine (and simple molecular substances generally) low melting and boiling points compared with giant structures. Because each iodine molecule is overall electrically neutral, and there are no ions or delocalised electrons free to move through the substance, iodine cannot carry an electric current, so it is a poor conductor.
In which structure is each carbon atom covalently bonded to four other carbon atoms?
Answer: A
Diamond's structure is a giant covalent lattice in which each carbon atom uses all four of its outer electrons to form four single covalent bonds to four other carbon atoms, building a rigid three-dimensional network. This is different from graphite, where each carbon atom bonds to only three others within flat layers.
Which property of graphite makes it useful as an electrode in electrolysis?
Answer: B
An electrode needs to conduct electricity and must not melt or break down under the conditions of electrolysis. Graphite conducts electricity because, within each layer, each carbon atom bonds to only three others, leaving one electron per atom delocalised and free to move along the layers. Graphite also has a high melting point because breaking its many strong covalent bonds takes a large amount of energy. Together, these properties make it a good, stable electrode material.
Describe the structure and bonding in diamond.
Describe the structure and bonding in graphite, and explain how this structure allows graphite to be used as a lubricant.
State one other use of graphite and one use of diamond, and for each one link the use to a property of its structure.
Diamond and graphite are both forms of carbon but have very different structures. In diamond, every carbon atom forms four strong covalent bonds to four neighbours, producing a rigid three-dimensional lattice with no weak points - this rigidity is why diamond is extremely hard and is used in cutting tools. In graphite, each carbon atom bonds to only three neighbours, forming flat layers of hexagonal rings; the covalent bonds within each layer are strong, but the forces between separate layers are weak, so the layers can slide over one another - this is why graphite feels slippery and is used as a lubricant, and also, because one electron per carbon atom is left delocalised and free to move along the layers, why graphite conducts electricity and can be used as an electrode.
Silicon(IV) oxide, SiO2, has a giant covalent structure similar to diamond.
Describe the structure of silicon(IV) oxide.
Explain why silicon(IV) oxide, like diamond, has a very high melting point and does not conduct electricity.
Silicon(IV) oxide has a giant covalent structure much like diamond: instead of individual molecules, every atom is linked to its neighbours by strong covalent bonds that extend throughout the whole structure (each silicon atom bonded to four oxygens, each oxygen bridging two silicons). Melting this structure means breaking a huge number of these strong covalent bonds at once, which needs a very large amount of energy - this is why silicon(IV) oxide, like diamond, has such a high melting point. It does not conduct electricity because, unlike a metal or a graphite layer, it has no delocalised electrons free to move, and unlike an ionic compound, it contains no ions - there is simply nothing present that is free to carry an electric charge.
Which description best fits the structure of a metal such as copper?
Answer: C
In a metal, each metal atom loses its outer electrons to the whole structure; these electrons become delocalised, meaning they are no longer attached to any one atom and are free to move throughout the metal. What is left behind is a regular lattice of positive metal ions, held together by their electrostatic attraction to this surrounding "sea" of delocalised electrons.
Describe the structure of metallic bonding in a metal such as copper.
Explain, in terms of structure and bonding, why metals are good conductors of electricity.
Explain, in terms of structure and bonding, why metals can be bent and shaped without breaking.
A metal is a giant lattice of positive metal ions surrounded by a sea of delocalised electrons - electrons that have left their original atoms and are free to move anywhere throughout the structure. This mobility explains both key properties. Electrical conductivity: because the delocalised electrons can move freely, they can flow through the metal and carry an electric current when a voltage is applied. Malleability and ductility: when a force is applied, the layers of positive ions can slide past one another into new positions; because the delocalised electrons are not tied to any particular ion, they simply move along too and continue to hold the ions together in their new arrangement. This is why metals can be hammered into sheets or drawn into wires without breaking, unlike ionic compounds, where sliding layers brings ions of the same charge together, causing them to repel and the structure to shatter.