Assessed in Paper 1A (multiple choice), Paper 1B (data-based) and Paper 2 (short and extended response). IB Chemistry guide (first assessment 2025).
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In preparation: 0 of 3 sections compiled and reviewed. The rest show key terms and common misconceptions from our question bank until they are.
S1.2.1 Nucleus
Nucleus
The small, dense, positively charged centre of an atom, made of protons and neutrons. It contains nearly all of the atom's mass but has a diameter roughly 10 000 to 100 000 times smaller than that of the atom (of order 10⁻¹⁵ m compared with 10⁻¹⁰ m). Its charge is positive because it contains the protons; the neutrons have no charge.
Nucleon
A particle found in the nucleus: a proton or a neutron. The mass number A of an atom is the number of nucleons it contains.
Proton
A nucleon with relative mass 1 and relative charge +1. The number of protons in the nucleus is the atomic number Z, and it fixes which element the atom belongs to.
Neutron
A nucleon with relative mass 1 and relative charge 0. Neutrons add to the mass of the nucleus but not to its charge. Number of neutrons = A − Z.
Electron
A particle with relative charge −1 and relative mass about 0.0005 (about 1/1836 of the mass of a proton), so its mass can be considered negligible. Electrons occupy the space outside the nucleus. A neutral atom has equal numbers of electrons and protons; the charge of an electron is equal in size and opposite in sign to the charge of a proton.
Atomic number (Z)
The number of protons in the nucleus of an atom. All atoms of an element have the same atomic number, and in a neutral atom it is also the number of electrons. It has no units.
Mass number (A)
The number of protons plus neutrons (nucleons) in the nucleus of an atom. It has no units. Electrons are not counted, because they are not nucleons and their mass is negligible.
Nuclear symbol ᴬ_Z X
The notation for a particular nucleus: the element symbol X with the mass number A as a left superscript and the atomic number Z as a left subscript. For example, ³⁷₁₇Cl has 17 protons and 37 − 17 = 20 neutrons; a neutral atom has 17 electrons. An ion's charge is written as a right superscript, e.g. ³⁷₁₇Cl⁻.
Ion (particle counts)
An ion forms when an atom gains or loses electrons; the numbers of protons and neutrons do not change. Number of electrons = Z − (charge). For ⁵⁶₂₆Fe³⁺: 26 protons, 30 neutrons, 26 − 3 = 23 electrons. For ³³₁₆S²⁻: 16 protons, 17 neutrons, 16 + 2 = 18 electrons.
Students often think The number of electrons is the mass number minus the atomic number (A − Z). In fact In a neutral atom the number of electrons equals the atomic number Z. The subtraction A − Z gives the number of neutrons, not electrons.
Students often think A positive ion has gained protons (and a negative ion has lost protons): the charge on an ion comes from a change in the number of protons. In fact No. Ions form by the loss or gain of electrons. The number of protons (and neutrons) is unchanged, so the atomic number and the element stay the same.
S1.2.2 Isotopes
Isotopes
Atoms of the same element (same atomic number, so the same number of protons) with different numbers of neutrons, and therefore different mass numbers. Neutral atoms of isotopes have the same number and arrangement of electrons.
Relative isotopic mass
The mass of an atom of one isotope relative to one-twelfth of the mass of an atom of carbon-12. It has no units and is close to, but not exactly equal to, the mass number (e.g. 34.97 for chlorine-35).
Percentage abundance
The percentage of the atoms of an element, in a sample, that are atoms of a particular isotope. The percentage abundances of all the isotopes of an element add up to 100%.
Relative atomic mass (A_r)
The weighted mean of the relative isotopic masses of an element's isotopes, each weighted by its abundance: A_r = Σ(relative isotopic mass × percentage abundance) ÷ 100. It has no units and is usually not a whole number, because most elements are mixtures of isotopes. It lies between the lightest and heaviest isotopic masses. For an element with two isotopes it lies closer to the mass of the more abundant isotope.
Physical and chemical properties of isotopes
Isotopes of an element have the same number and arrangement of electrons, so they undergo the same chemical reactions. Their atoms differ in mass, so physical properties that depend on mass differ: for example density, rate of diffusion, and melting and boiling points (D₂O, water made from hydrogen-2, melts at 3.8 °C and boils at 101.4 °C).
Students often think Atoms with different mass numbers are atoms of different elements, because their masses differ. In fact Yes. The element is defined by the atomic number (number of protons). Atoms of one element with different mass numbers are isotopes.
Students often think Isotopes of an element have different chemical properties, because their atoms have different masses (and some isotopes are radioactive). In fact No. Chemical reactions involve electrons, and isotopes have the same number and arrangement of electrons, so they undergo the same chemical reactions.
S1.2.3 Mass spectrum of an element HL
Mass spectrum of an element
A plot of relative abundance against m/z for the ions formed from a sample. For an element, each peak from singly charged atomic ions corresponds to one isotope: its m/z value gives the isotope's relative mass and its height gives the isotope's relative abundance. The number of such peaks is the number of isotopes present.
Mass-to-charge ratio (m/z)
The mass of an ion divided by its charge number. For a singly charged ion (z = 1), m/z equals the relative mass of the ion, which is the relative isotopic mass of the atom (the mass of the lost electron being negligible).
Relative abundance (in a mass spectrum)
The height of a peak, often scaled so that the tallest peak is 100. Relative abundances scaled in this way are not percentages: A_r = Σ(m/z × relative abundance) ÷ Σ(relative abundances).
Students often think The relative atomic mass of an element is the m/z value of the tallest peak in its mass spectrum. In fact No. The tallest peak gives the mass of the most abundant isotope. A_r is the abundance-weighted mean of the m/z values of all the peaks.
Students often think The relative abundances shown in a mass spectrum are percentages, so Σ(m/z × relative abundance) is always divided by 100. In fact Not necessarily. When the tallest peak is scaled to 100, the relative abundances add up to more than 100, so the weighted sum must be divided by their total, not by 100.
Diagnostic a bearings check, not a test
6 questions, one per part of the topic where we can. Answer them, then see which statements you own and which to read.
1 Which statement about the nucleus of an atom is correct?
Answer and reasoning
It takes up most of the atom's volume, with electrons at its surface. — Textbook diagrams draw the nucleus as a large ball, but they are not to scale. The nucleus is roughly 10 000 to 100 000 times smaller in diameter than the atom; the electrons occupy almost all of the atom's volume.
Its density is far greater than the average density of the atom. — Correct. Nearly all of the atom's mass is in the nucleus, because electrons are so light. The nucleus is of order 10⁻¹⁵ m across and the atom of order 10⁻¹⁰ m, so the nucleus occupies only about (10⁻¹⁵ ÷ 10⁻¹⁰)³ = 10⁻¹⁵ of the atom's volume. Nearly all the mass in a tiny volume makes the nucleus extremely dense.
It contains only about two-thirds of the mass of the atom. — If electrons were as heavy as protons, an atom with similar numbers of protons, neutrons and electrons would carry about a third of its mass outside the nucleus. In fact an electron's relative mass is about 0.0005, compared with 1 for a proton or neutron, so electrons contribute a negligible fraction and nearly all of the mass is in the nucleus.
It contains the electrons as well as the protons and neutrons. — The nucleus contains only nucleons: protons and neutrons. The electrons occupy the space outside the nucleus.
2 Two atoms have the same atomic number but different mass numbers. Which statement about these atoms is correct?
Answer and reasoning
They belong to different elements, as they contain different numbers of nucleons in total. — The element is fixed by the number of protons, not by the number of nucleons. The same atomic number means the same element; the extra nucleons are neutrons.
They belong to the same element, as they differ only in their numbers of neutrons. — Correct. The atomic number (number of protons) defines the element. Different mass numbers with the same atomic number mean different numbers of neutrons: the atoms are isotopes of one element.
They have different numbers of electrons, as their mass numbers are different. — Neutral atoms have as many electrons as protons, and both atoms have the same atomic number, so they have the same number of electrons. A different number of electrons would make an ion, not an isotope.
They undergo different chemical reactions, because their atoms have different masses. — Chemical reactions depend on electrons, and isotopes have the same number and arrangement of electrons, so they undergo the same reactions. Only mass-dependent physical properties differ.
3 The mass spectrum of a sample of copper shows two peaks from singly charged atomic ions: m/z 63 (relative abundance 100) and m/z 65 (relative abundance 44.6). What do these peaks show? HL
Answer and reasoning
The sample also contains a second element, of mass 65. — Atoms of one element can have different masses: these are isotopes. Different mass numbers do not mean different elements; the element is fixed by the number of protons.
The A_r of copper is 63, which is the m/z of the tallest peak. — The tallest peak gives the mass of the most abundant isotope, copper-63. A_r is the weighted mean of both peaks: (63 × 100 + 65 × 44.6) ÷ 144.6 = 63.62 from these whole-number values.
The A_r of copper is 64, the mean of the two m/z values. — 64 is the simple mean, correct only if both isotopes were equally abundant. The peak at 63 is more than twice as tall, so A_r lies closer to 63.
Copper atoms exist with two numbers of neutrons, 34 and 36. — Correct. Each peak from singly charged atomic ions corresponds to one isotope. Copper atoms all have the same number of protons, so masses of 63 and 65 mean two different numbers of neutrons (34 and 36).
4 Which statement about the relative masses and relative charges of subatomic particles is correct?
Answer and reasoning
A proton carries a larger charge than an electron does. — Mass and charge are separate properties. The proton is about 1836 times heavier than the electron, but their charges are equal in size (+1 and −1), which is why a neutral atom has equal numbers of protons and electrons.
A neutron carries a charge opposite in sign to a proton's. — The neutron's relative charge is 0: it carries no charge. The positive charge of the protons is balanced by the electrons outside the nucleus, not by the neutrons.
A proton and an electron carry charges that are equal in size. — Correct. Relative charges: proton +1, electron −1, equal in size and opposite in sign. This is why a neutral atom has equal numbers of protons and electrons. (Their masses are very different: proton 1, electron about 0.0005.)
An electron has a mass close to that of a proton. — The electron's relative mass is about 0.0005 (about 1/1836 of a proton's), so its mass can be considered negligible. Only its charge matches the proton's in size.
5 Chlorine (atomic number 17) has isotopes with mass numbers 35 and 37. Pure samples of ³⁵Cl₂ gas and ³⁷Cl₂ gas are compared at the same temperature and pressure. Which statement is correct?
Answer and reasoning
The ³⁷Cl₂ gas diffuses more slowly than ³⁵Cl₂. — Correct. At the same temperature the molecules of both gases have the same mean kinetic energy, so the heavier ³⁷Cl₂ molecules move more slowly on average and the gas diffuses more slowly. Rate of diffusion is a physical property that differs between isotopes.
The two gases diffuse at the same rate as each other. — Isotopes have the same chemical properties, but not the same physical properties. ³⁷Cl₂ molecules (relative mass about 74) are heavier than ³⁵Cl₂ molecules (about 70), so at the same temperature they move more slowly and diffuse more slowly.
The two gases undergo different chemical reactions. — Both isotopes have 17 electrons with the same arrangement, so ³⁵Cl₂ and ³⁷Cl₂ undergo the same chemical reactions. The difference between them is in mass-dependent physical properties such as rate of diffusion.
The ³⁷Cl₂ molecules contain more electrons than ³⁵Cl₂. — Every chlorine atom has 17 electrons whatever its mass number, so both molecules contain 34 electrons. The isotopes differ in their numbers of neutrons (18 and 20).
6 The mass spectrum of element E shows three peaks from singly charged atomic ions: m/z 24 (relative abundance 100), m/z 25 (relative abundance 12.7) and m/z 26 (relative abundance 13.9). What is the relative atomic mass of E, calculated from these data? HL
Answer and reasoning
24.32 — Correct. The relative abundances add up to 126.6, not 100. A_r = (24 × 100 + 25 × 12.7 + 26 × 13.9) ÷ 126.6 = 3078.9 ÷ 126.6 = 24.32.
30.79 — Dividing the weighted sum 3078.9 by 100 treats the relative abundances as percentages, but they add up to 126.6. A_r must lie between 24 and 26; 30.79 is outside that range.
24.00 — 24 is the m/z of the tallest peak, the most abundant isotope. The other two isotopes also contribute, so A_r is a weighted mean that lies a little above 24.
25.00 — 25 is the simple mean of 24, 25 and 26, which would be correct only if the three isotopes were equally abundant. The isotope with m/z 24 is far more abundant, so each value must be weighted.
Read the ones marked not yet in Learn, then Verify.
Verify confirm before you go
8 more questions. Every wrong answer here is a real misconception, and you see why it is wrong straight away.
1 An atom of sodium, ²³₁₁Na (mass number 23, atomic number 11), forms the ion Na⁺. How does the mass of the ion compare with the mass of the atom?
Answer and reasoning
It is smaller by a negligible amount, because one electron was lost. — Correct. A 1+ ion has one electron fewer than its atom; the numbers of protons and neutrons do not change. An electron's relative mass is about 0.0005, so the loss is negligible (about 0.002% of the mass of the atom).
It is larger by a negligible amount, because one electron has been gained. — A positive ion has LOST an electron: removing a negatively charged particle leaves 11 protons and only 10 electrons, giving a charge of 1+. Gaining an electron would give a negative ion.
It is larger by about one unit, because the ion has gained one proton. — Ions form by the transfer of electrons only. Gaining a proton would change the atomic number from 11 to 12 and make the particle an ion of a different element (magnesium).
It is smaller by about one unit, because the ion has lost one electron. — The ion has indeed lost one electron, but an electron's relative mass is about 0.0005, not 1. The mass decreases by a negligible amount; the mass number stays 23.
2 How many electrons are there in one ³³₁₆S²⁻ ion (mass number 33, atomic number 16, charge 2−)?
Answer and reasoning
14 — A negative ion has GAINED electrons, because electrons carry negative charge. Subtracting the 2 gives 14, which would be the ion S²⁺. The correct count is 16 + 2 = 18.
19 — Taking A − Z = 33 − 16 = 17 as the atom's electrons and adding 2 for the charge gives 19. But 17 is the number of neutrons, not electrons. Electrons are not nucleons and are not counted in A. Electrons = Z − (charge) = 16 − (−2) = 18.
35 — This treats 33 as the atomic number. In ᴬ_Z X the superscript (33) is the mass number and the subscript (16) is the atomic number, so the ion has 16 protons and 16 + 2 = 18 electrons.
18 — Correct. Z = 16, so there are 16 protons. A charge of 2− means two more electrons than protons: 16 + 2 = 18 electrons. (The ion has 33 − 16 = 17 neutrons.)
Working Protons = Z = 16. Electrons = Z − (charge) = 16 − (−2) = 18. (Neutrons = A − Z = 33 − 16 = 17.)
3 An ion X³⁺ contains 23 electrons and 30 neutrons. Which nuclear symbol represents this ion? (Each symbol is written with its superscript first and its subscript second.)
Answer and reasoning
⁵⁰₂₀X³⁺ — This assumes the 3+ ion has GAINED three electrons, giving 23 − 3 = 20 protons. A positive ion has LOST electrons, so it has 23 + 3 = 26 protons, and A = 26 + 30 = 56.
⁵³₂₃X³⁺ — This sets the number of protons equal to the number of electrons, as in a neutral atom. The 3+ charge means there are three more protons than electrons: Z = 26 and A = 56.
⁵⁶₂₆X³⁺ — Correct. A 3+ ion has three fewer electrons than protons, so protons = 23 + 3 = 26 = Z. Mass number A = protons + neutrons = 26 + 30 = 56.
²⁶₅₆X³⁺ — The counts are right (26 protons, 56 nucleons) but the numbers are in the wrong places. The mass number is the superscript and the atomic number is the subscript; A cannot be smaller than Z.
Working Charge = protons − electrons, so protons = 23 + 3 = 26 = Z. A = protons + neutrons = 26 + 30 = 56. Symbol ⁵⁶₂₆X³⁺ (the element is iron).
4 Hydrogen (atomic number 1) has isotopes ¹H and ²H. Water made from ²H, D₂O, melts at 3.8 °C, boils at 101.4 °C and has a density of about 1.1 g cm⁻³; ordinary water, H₂O, melts at 0.0 °C, boils at 100.0 °C and has a density of about 1.0 g cm⁻³. Both react with sodium in the same way: H₂O forms NaOH and H₂, and D₂O forms NaOD and D₂. Which statement accounts for these observations?
Answer and reasoning
D₂O molecules contain more electrons, so the forces between them are stronger. — Both isotopes of hydrogen have atomic number 1, so a neutral atom of either has one electron, and H₂O and D₂O molecules both have 10 electrons. ²H differs from ¹H by one neutron; a change in the number of electrons would give an ion, not an isotope.
Boiling D₂O breaks its O–D bonds, which need more energy to break than O–H bonds. — Boiling a molecular substance separates whole molecules by overcoming the intermolecular forces between them; no O–D or O–H covalent bonds break. So the strength of the bonds inside the molecules does not set the boiling point, and this statement also leaves the melting points and densities unexplained.
The extra neutrons in D₂O carry charge, so D₂O molecules attract one another more. — Neutrons have no charge, so the extra neutrons do not change the charge of the nucleus or the atom. They add mass only, and it is the difference in mass that changes the physical properties.
D₂O molecules are heavier than H₂O molecules, but have the same electron arrangement. — Correct. ¹H has 1 proton and no neutron; ²H has 1 proton and 1 neutron. The electron arrangement, which controls chemical reactions, is the same, so the reactions are the same. The greater mass changes mass-dependent physical properties: melting point, boiling point and density.
5 A sample of chlorine contains 75.8% chlorine-35 (relative isotopic mass 34.97) and 24.2% chlorine-37 (relative isotopic mass 36.97). What is the relative atomic mass of chlorine in this sample?
Answer and reasoning
35.97 — This is the simple mean (34.97 + 36.97) ÷ 2, which would be correct only if the isotopes were equally abundant. Chlorine-35 is about three times as abundant, so each mass must be weighted by its abundance.
35.45 — Correct. A_r = (75.8 × 34.97 + 24.2 × 36.97) ÷ 100 = (2650.7 + 894.7) ÷ 100 = 35.45. It lies closer to 34.97 because chlorine-35 is the more abundant isotope.
34.97 — This is the relative isotopic mass of the most abundant isotope only. Chlorine-37 makes up almost a quarter of the atoms and raises the weighted mean to 35.45.
17.73 — The weighted sum 0.758 × 34.97 + 0.242 × 36.97 = 35.45 is already the mean; dividing it by 2 (the number of isotopes) is an extra step. A_r must lie between 34.97 and 36.97.
6 Copper has two isotopes, copper-63 and copper-65, with relative isotopic masses 62.93 and 64.93. The relative atomic mass of copper is 63.55. What is the percentage abundance of copper-65?
Answer and reasoning
69.0% — (64.93 − 63.55) ÷ 2.00 = 0.690 is the fraction of copper-63, not copper-65. A_r lies closer to the more abundant isotope; 63.55 is closer to 62.93, so copper-65 must be the minor isotope.
49.7% — Writing both abundances as x (62.93x + 64.93x = 6355) ignores the fact that the abundances add up to 100%. If copper-65 is x%, copper-63 is (100 − x)%.
27.5% — Using the mass numbers 63 and 65 instead of the isotopic masses 62.93 and 64.93 gives (63.55 − 63) ÷ 2 = 0.275. The stem gives the relative isotopic masses; with them the fraction is 0.62 ÷ 2.00 = 0.310.
31.0% — Correct. Let x% be copper-65: 64.93x + 62.93(100 − x) = 6355, so 2.00x = 62, x = 31.0. A check: A_r (63.55) is closer to 62.93, so copper-63 should be the more abundant isotope, as it is (69.0%).
Working Let the abundance of copper-65 be x%. Then 64.93x + 62.93(100 − x) = 63.55 × 100 → 6293 + 2.00x = 6355 → x = 62 ÷ 2.00 = 31.0%. (Copper-63: 69.0%.)
7 The mass spectrum of element X, showing only the peaks from singly charged atomic ions (X⁺), has two peaks: m/z 79 (relative abundance 100) and m/z 81 (relative abundance 97.3). Atomic numbers: selenium 34, bromine 35, gadolinium 64, gold 79. Which element is X? HL
Answer and reasoning
Element X is selenium (A_r 78.97). — This takes A_r as 79, the m/z of the tallest peak. The peak at m/z 81 is almost as tall, so the weighted mean is about 80, which matches bromine.
Element X is gold (A_r 196.97). — This reads m/z 79 as the atomic number. m/z of a singly charged ion is its relative mass, which is related to the mass number, not the atomic number. Gold atoms have masses near 197.
Element X is bromine (A_r 79.90). — Correct. A_r = (79 × 100 + 81 × 97.3) ÷ (100 + 97.3) = 15 781.3 ÷ 197.3 = 79.99, which matches bromine (79.90) and no other option. The two peaks are two isotopes of nearly equal abundance.
Element X is gadolinium (A_r 157.25). — Dividing 15 781.3 by 100 treats the relative abundances as percentages, but they add up to 197.3. The weighted mean of the two m/z values must lie between 79 and 81.
Working Σ(relative abundance) = 100 + 97.3 = 197.3. A_r = (79 × 100 + 81 × 97.3) ÷ 197.3 = (7900 + 7881.3) ÷ 197.3 = 79.99 → closest to bromine (A_r 79.90). (m/z values are whole numbers here, so the calculated value differs slightly from the accepted A_r.)
8 How many neutrons are there in one atom of ⁸¹₃₅Br (mass number 81, atomic number 35)?
Answer and reasoning
35 — 35 is the number of protons (and of electrons in the neutral atom). The number of neutrons is not tied to the atomic number: it is A − Z = 81 − 35 = 46.
46 — Correct. The mass number 81 counts the protons and neutrons; the atomic number 35 is the number of protons. Neutrons = A − Z = 81 − 35 = 46.
81 — 81 is the mass number, the total number of protons and neutrons. Subtract the 35 protons to get the neutrons: 81 − 35 = 46.
11 — Subtracting 35 electrons as well as 35 protons treats the electrons as if they had a mass like a proton's and were counted in A. Electrons are not nucleons and their mass is negligible, so A counts only protons and neutrons: 81 − 35 = 46.
Working Protons = Z = 35. Mass number A = protons + neutrons, so neutrons = A − Z = 81 − 35 = 46.
That was your twenty minutes. Real practice on S1.2 is past-paper questions marked against the mark scheme.
Paper 1A tests it as multiple choice; Paper 1B through data you have not seen before; Paper 2 with short answers and, at HL, extended responses. Look for the command words — outline, explain, compare, evaluate — and give exactly what each asks for.
Compiled from the IB Chemistry guide (first assessment 2025) and our question bank · Specialist review in progress. How these pages are made · Free, no account ·
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