DP Chemistry HL S1.2 The Nuclear Atom

DP Chemistry HL S1.2 The Nuclear Atom - Full notes

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DP Chemistry HL | S1.2 The Nuclear Atom

DP CHEMISTRY · STRUCTURE

S1.2

The Nuclear Atom

Full notes for HL · IB-style

Contents

1	The structure of the atom
2	Atomic number, mass number and nuclear symbols
3	Ions
4	Isotopes
5	Relative atomic mass
6	Mass spectrometry
7	The development of atomic models
8	Subatomic particles: properties
9	Nuclear notation and counting particles
10	Isotopes and their properties
11	Relative atomic mass from isotopic abundance
12	Finding abundances from relative atomic mass
13	Mass spectrometry
14	Mass spectra of diatomic elements
15	Radioisotopes and their uses
16	Half-life calculations
17	Ions: formation and notation
18	Precision in definitions
19	Calculations with masses of particles
20	Mixed review: worked problems
21	Additional exam-style practice
22	Common misconceptions and links
23	Exam-style question with markscheme
24	Check your understanding
25	Summary and topic map

These are full teaching notes. Read each section in order, try the Think about it question before you read on, and do every Quick check before moving to the next section.

Before you start

Rate yourself on each statement now, and again after you've worked through the notes. Shade R (not yet), A (getting there) or G (confident).

I can…	R	A	G
describe the nuclear model of the atom and the properties of subatomic particles
use nuclear symbols to deduce numbers of protons, neutrons and electrons in atoms and ions
define isotopes and explain their similar chemical and different physical properties
calculate relative atomic mass from isotopic abundances
interpret mass spectra to find isotopic composition

There is no additional HL content in this subtopic. SL and HL students study the same material.

Guiding questions

How can we model the structure of the atom? How do isotopes explain non-integer relative atomic masses?

How this topic is examined

S1.2 is examined through nuclear symbol questions, relative atomic mass calculations and mass spectra. Expect at least one calculation of Ar from abundance data in most papers.

Give calculated relative atomic masses to two decimal places and remember that Ar has no units.

Command term	What the examiner wants
State / Identify	a brief answer, no explanation
Outline / Describe	a brief or detailed account, without reasons
Explain	a detailed account with causes, reasons or mechanisms
Calculate / Determine	a numerical answer with working, correct significant figures and units
Deduce / Predict	reach a conclusion from given information
Draw / Sketch	a clear labelled diagram, structure or graph
Evaluate / Discuss	weigh strengths and limitations; reach a judgement

How to use these notes

Each section teaches one part of the syllabus in full. You'll see these boxes throughout; each has one job:

Box	What it's for
Think about it	a question to get you thinking before a section; the answer comes in the section
Key point	the one idea you must not forget
Common mistake	an error examiners see every year, and how to avoid it
In the real world	where the chemistry shows up outside the classroom
Try this	something quick you can do yourself to test the idea
Worked example	full solutions with IB-style marks (M1 method, A1 answer, R1 reasoning)
Quick check	short questions to test yourself at the end of each section

1 The structure of the atom

A tiny dense nucleus surrounded by electrons.

Think about it

If an atom were the size of a football stadium, the nucleus would be about the size of a pea at the centre. Where is almost all the mass of the atom?

Atoms consist of a very small, dense, positively charged nucleus, containing protons and neutrons (together called nucleons), surrounded by electrons moving in the much larger space around it. The nucleus contains almost all the mass of the atom but occupies a tiny fraction of its volume: the radius of the nucleus is about 10−15 m, while that of the atom is about 10−10 m.

This nuclear model came from Rutherford's gold foil experiment (1909, carried out by Geiger and Marsden). Alpha particles fired at thin gold foil mostly passed straight through, showing that the atom is mostly empty space; a few were deflected at large angles, and a very few bounced back, showing a small, dense, positive nucleus.

Atoms are electrically neutral because the number of protons equals the number of electrons.

Particle	Relative mass	Relative charge	Location
Proton	1	+1	nucleus
Neutron	1	0	nucleus
Electron	1/1836 (about 0.0005)	−1	around the nucleus

Quick check

1. Which subatomic particles are in the nucleus?

2. What did large-angle deflections in Rutherford's experiment show?

Answers: 1. protons and neutrons 2. a small, dense, positive nucleus

2 Atomic number, mass number and nuclear symbols

Two numbers identify an atom.

The atomic number (Z) is the number of protons in the nucleus. It defines the element: every atom of carbon has 6 protons. In a neutral atom it also equals the number of electrons.

The mass number (A) is the total number of protons and neutrons in the nucleus. The number of neutrons = A − Z.

These are shown in nuclear symbol notation, with A as a superscript and Z as a subscript before the element symbol, for example 2311Na: sodium-23 has 11 protons, 12 neutrons and 11 electrons.

WORKED EXAMPLE 1 Using nuclear symbols

Deduce the numbers of protons, neutrons and electrons in 3717Cl and in 5626Fe3+.

Solution

Cl: 17 protons, 37 − 17 = 20 neutrons, 17 electronsA1

Fe3+: 26 protons, 56 − 26 = 30 neutrons, 26 − 3 = 23 electronsA1

Quick check

1. Define mass number.

2. How many neutrons are in 146C?

Answers: 1. number of protons plus neutrons 2. 8

3 Ions

Atoms gain or lose electrons to form charged particles.

An ion is a charged particle formed when an atom (or group of atoms) gains or loses electrons. The number of protons does not change. Cations are positive ions, formed by losing electrons; metals typically form cations (Na+, Mg2+). Anions are negative ions, formed by gaining electrons; non-metals typically form anions (Cl−, O2−).

Number of electrons in an ion = Z − charge. For example, O2− has 8 protons and 8 + 2 = 10 electrons.

Species	Protons	Neutrons	Electrons
24Mg2+	12	12	10
19F−	9	10	10
40Ca	20	20	20
32S2−	16	16	18

Common mistake

Changing the number of protons when forming an ion. Only electrons are gained or lost; the proton number defines the element and stays the same.

4 Isotopes

Same element, different numbers of neutrons.

Isotopes are atoms of the same element (same atomic number, same number of protons) with different numbers of neutrons, and so different mass numbers. Carbon has three naturally occurring isotopes: carbon-12, carbon-13 and carbon-14 (12C, 13C, 14C).

Isotopes have the same chemical properties, because chemical behaviour depends on electron arrangement, which is the same. They have different physical properties that depend on mass: density, rate of diffusion, and melting and boiling points (slightly). Some isotopes are radioactive (radioisotopes) because their nuclei are unstable.

Uses of radioisotopes: carbon-14 for radiocarbon dating of organic material; cobalt-60 in radiotherapy for cancer; iodine-131 as a medical tracer and treatment for thyroid conditions; technetium-99m in medical imaging.

Quick check

1. Define isotopes.

2. Why do isotopes have the same chemical properties?

Answers: 1. atoms of the same element with different numbers of neutrons 2. same number and arrangement of electrons

5 Relative atomic mass

A weighted average of isotope masses.

Masses of atoms are compared with carbon-12, which is defined as having a mass of exactly 12 atomic mass units. The relative atomic mass (Ar) of an element is the weighted average mass of its atoms, taking into account the natural abundance of each isotope, compared with 1/12 of the mass of a carbon-12 atom. It has no units.

Because most elements are mixtures of isotopes, Ar values are usually not whole numbers: chlorine is 35.45, because it contains about 75.8% chlorine-35 and 24.2% chlorine-37.

Ar = Σ (isotopic mass × percentage abundance) ÷ 100

WORKED EXAMPLE 2 Calculating relative atomic mass

Chlorine consists of 75.78% 35Cl and 24.22% 37Cl. Calculate Ar.

Solution

Ar = (35 × 75.78 + 37 × 24.22) ÷ 100M1

= (2652.3 + 896.1) ÷ 100 = 35.48A1

WORKED EXAMPLE 3 Finding an abundance

Boron has Ar = 10.81 and consists of 10B and 11B. Calculate the percentage abundance of 10B.

Solution

Let x = % of 10B: 10x + 11(100 − x) = 1081M1

1100 − x = 1081, so x = 19%A1

Quick check

1. Why is Ar for chlorine not a whole number?

2. Gallium is 60.1% 69Ga and 39.9% 71Ga. Calculate Ar.

Answers: 1. it is a weighted average of isotopes 2. 69.80

6 Mass spectrometry

Measuring the masses and abundances of isotopes.

A mass spectrometer measures the masses and relative abundances of particles. In outline: the sample is vaporized; atoms are ionized to form positive ions (by bombarding with electrons); ions are accelerated by an electric field; they are deflected or separated according to their mass to charge ratio (m/z); and they are detected. Lighter ions (or ions of higher charge) are deflected more.

The result is a mass spectrum: a graph of relative abundance against m/z. For an element, each peak corresponds to an isotope (for ions with charge +1, m/z equals the mass). The heights of the peaks give the relative abundances, from which Ar is calculated. (Mass spectra of compounds, which show fragments, are covered in S3.2.)

Figure 1. Mass spectrum of zirconium (m/z values on the x-axis).

WORKED EXAMPLE 4 Ar from a mass spectrum

Using the spectrum of zirconium, calculate Ar.

Solution

Ar = (90 × 51.5 + 91 × 11.2 + 92 × 17.1 + 94 × 17.4 + 96 × 2.8) ÷ 100M1

= 9131.8 ÷ 100 = 91.32A1

Quick check

1. What does the position of a peak show?

2. What does the height of a peak show?

Answers: 1. m/z, the mass of the isotope (for +1 ions) 2. relative abundance

7 The development of atomic models

From Dalton to Rutherford and Bohr.

Dalton (1803) proposed indivisible atoms of each element. Thomson (1897) discovered the electron and proposed the 'plum pudding' model. Rutherford's gold-foil experiment (1911) showed most of the atom is empty space with a tiny, dense, positive nucleus. Bohr (1913) placed electrons in fixed energy levels. Chadwick (1932) discovered the neutron.

Each model was replaced when new evidence could not be explained: this is how scientific models develop.

WORKED EXAMPLE 5 Gold-foil evidence (standard)

In Rutherford's experiment, most alpha particles passed straight through gold foil, but a few were deflected through large angles. Explain what each observation shows.

Solution

Most passed through: the atom is mostly empty spaceR1

A few deflected strongly: a small, dense, positively charged nucleusR1

In the real world

Modern scanning tunnelling microscopes can image individual atoms on a surface, giving direct visual evidence for atoms that Dalton could only infer.

Quick check

1. Why did the discovery of the electron disprove Dalton's model?

2. What did Chadwick's discovery explain?

Answers: 1. atoms contain smaller particles, so they are not indivisible 2. the mass of atoms is greater than the mass of their protons alone

8 Subatomic particles: properties

Relative mass, charge and location.

The nucleus contains almost all the mass of the atom but occupies a tiny fraction of its volume: the nucleus is about 10−15 m across, the atom about 10−10 m. Protons and neutrons are collectively called nucleons.

Particle	Relative mass	Relative charge	Location
proton	1	+1	nucleus
neutron	1	0	nucleus
electron	1/1836 (≈ 0.0005)	−1	energy levels around the nucleus

WORKED EXAMPLE 6 Scale (standard)

If a nucleus were the size of a football (22 cm), estimate the diameter of the atom.

Solution

Ratio 10−10/10−15 = 105M1

22 cm × 105 = 22 kmA1

Quick check

1. Which particle has negligible mass?

2. Why is an atom electrically neutral?

Answers: 1. electron 2. it has equal numbers of protons and electrons

9 Nuclear notation and counting particles

A, Z and charge.

The notation AZX gives the mass number A (protons + neutrons) and the atomic number Z (protons). Number of neutrons = A − Z. For ions, number of electrons = Z − charge.

Species	Protons	Neutrons	Electrons
3517Cl	17	18	17
3717Cl−	17	20	18
5626Fe3+	26	30	23
3216S2−	16	16	18

WORKED EXAMPLE 7 Counting (basic)

State the numbers of protons, neutrons and electrons in 2713Al3+.

Solution

13 protons, 14 neutronsA1

13 − 3 = 10 electronsA1

WORKED EXAMPLE 8 Identifying (standard)

A species has 20 protons, 20 neutrons and 18 electrons. Write its symbol.

Solution

Z = 20: calcium; A = 40; charge +2M1

4020Ca2+A1

Quick check

1. Find the numbers of particles in 7935Br−.

2. Which two species in the table are isoelectronic (same number of electrons)?

Answers: 1. 35 p, 44 n, 36 e 2. Cl− and S2− (18 each)

10 Isotopes and their properties

Same element, different mass.

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have the same chemical properties (same electron configuration) but different physical properties such as mass, density, rate of diffusion and, for some, radioactivity.

WORKED EXAMPLE 9 Properties (standard)

Compare the chemical and physical properties of 1H2 and 2H2 (deuterium).

Solution

Same chemical reactions: same electron arrangementR1

D2 is denser and diffuses more slowly: different physical propertiesR1

In the real world

Heavy water (D2O) is used as a moderator in some nuclear reactors. It has a higher boiling point (101.4 °C) and density than normal water, but reacts in the same way.

Quick check

1. Why do isotopes have the same chemical properties?

2. State one physical property that differs between 35Cl2 and 37Cl2.

Answers: 1. same number of electrons, so same electron configuration 2. density (or mass, rate of diffusion)

11 Relative atomic mass from isotopic abundance

A weighted mean.

Relative atomic mass Ar is the weighted mean mass of the atoms of an element relative to 1/12 of the mass of a carbon-12 atom. Ar = Σ(isotope mass × % abundance)/100. Ar has no units.

WORKED EXAMPLE 10 Two isotopes (standard)

Chlorine is 75.78% 35Cl (34.97) and 24.22% 37Cl (36.97). Calculate Ar.

Solution

(34.97 × 75.78 + 36.97 × 24.22)/100M1

= 35.45A1

WORKED EXAMPLE 11 Three isotopes (standard)

Magnesium: 24Mg 78.99%, 25Mg 10.00%, 26Mg 11.01%. Calculate Ar (use mass numbers).

Solution

(24 × 78.99 + 25 × 10.00 + 26 × 11.01)/100M1

= 24.32A1

Quick check

1. Copper has isotopes 63Cu (69.2%) and 65Cu (30.8%). Calculate Ar.

2. Silicon: 28Si 92.2%, 29Si 4.7%, 30Si 3.1%. Calculate Ar.

Answers: 1. 63.6 2. 28.1

12 Finding abundances from relative atomic mass

Working backwards.

With two isotopes, let the abundance of one be x%, so the other is (100 − x)%. Substitute into the Ar expression and solve.

WORKED EXAMPLE 12 Boron (exam level)

Boron has Ar = 10.81 and isotopes 10B and 11B. Calculate the percentage abundance of each.

Solution

10x + 11(100 − x) = 1081M1

x = 19A1

10B 19%, 11B 81%A1

Quick check

1. Gallium (Ar = 69.72) has isotopes 69Ga and 71Ga. Find their abundances.

2. Rubidium (Ar = 85.47) has isotopes 85Rb and 87Rb. Find their abundances.

Answers: 1. 69Ga 64%, 71Ga 36% 2. 85Rb 76.5%, 87Rb 23.5%

13 Mass spectrometry

Measuring isotopic masses and abundances.

A mass spectrometer ionizes atoms (or molecules), accelerates the ions, separates them according to their mass-to-charge ratio (m/z) and detects them. The mass spectrum of an element shows one peak for each isotope; the height of each peak is proportional to the abundance of that isotope.

Figure 2. Mass spectrum of magnesium (m/z on the horizontal axis).

WORKED EXAMPLE 13 Reading a spectrum (standard)

Use the mass spectrum of magnesium to calculate Ar.

Solution

(24 × 78.99 + 25 × 10.00 + 26 × 11.01)/100M1

= 24.32A1

Quick check

1. How many peaks would the mass spectrum of chlorine atoms show?

2. What does the height of a peak represent?

Answers: 1. 2 (m/z 35 and 37) 2. the relative abundance of that isotope

14 Mass spectra of diatomic elements

Molecular ions and combinations.

For a diatomic element such as Cl2, the molecular ion peaks come from all combinations of isotopes: 35Cl35Cl (m/z 70), 35Cl37Cl (72) and 37Cl37Cl (74). Their relative heights follow from multiplying the isotope fractions.

WORKED EXAMPLE 14 Chlorine molecules (exam level)

With 35Cl : 37Cl = 3 : 1, predict the ratio of the peaks at m/z 70, 72 and 74.

Solution

P(35,35) = 9/16; P(35,37) = 2 × 3/16 = 6/16; P(37,37) = 1/16M1 A1

Ratio 9 : 6 : 1A1

Quick check

1. Bromine has 79Br : 81Br ≈ 1 : 1. Predict the molecular ion peaks of Br2 and their ratio.

Answers: 1. m/z 158, 160, 162 in ratio 1 : 2 : 1

15 Radioisotopes and their uses

Unstable nuclei.

Some isotopes have unstable nuclei and are radioactive. Radioisotopes have many uses: 14C in radiocarbon dating; 60Co in radiotherapy and food irradiation; 131I as a medical tracer for the thyroid; 99mTc in medical imaging; 241Am in smoke detectors.

Radioisotope	Use	Why it is suitable
14C	dating once-living material	constant ratio in living things; known half-life (5730 years)
131I	thyroid diagnosis and treatment	taken up by the thyroid
60Co	radiotherapy	emits penetrating gamma radiation
241Am	smoke detectors	alpha particles ionize air; long half-life

In the real world

Radiocarbon dating can date organic materials up to about 50 000 years old, which has been crucial in archaeology, for example in dating the Dead Sea Scrolls.

Quick check

1. Why is 14C not useful for dating dinosaur fossils (65 million years old)?

2. Why must a medical tracer have a short half-life?

Answers: 1. almost all 14C would have decayed 2. to limit the radiation dose to the patient

16 Half-life calculations

Fraction remaining.

The half-life is the time for half the radioactive nuclei in a sample to decay. After n half-lives, the fraction remaining is (1/2)n.

WORKED EXAMPLE 15 Dating (standard)

A wooden artefact has 25% of the 14C of living wood. Estimate its age (half-life 5730 years).

Solution

25% = (1/2)2: 2 half-livesM1

Age = 11 460 yearsA1

WORKED EXAMPLE 16 Mass remaining (standard)

A 40 mg sample of 131I (half-life 8 days) is stored for 24 days. Find the mass remaining.

Solution

24/8 = 3 half-livesM1

40 × (1/2)3 = 5 mgA1

Quick check

1. What fraction of a radioisotope remains after 5 half-lives?

2. A sample's activity falls from 800 to 50 counts per minute in 20 hours. Find the half-life.

Answers: 1. 1/32 2. 4 half-lives: 5 hours

17 Ions: formation and notation

Losing and gaining electrons.

Metals tend to lose electrons to form positive ions (cations); non-metals tend to gain electrons to form negative ions (anions). The charge equals the number of electrons lost or gained. Many ions have the same electron arrangement as a noble gas.

WORKED EXAMPLE 17 Ion formation (standard)

Write the symbols and electron numbers for the ions formed by Mg and O.

Solution

Mg → Mg2+ + 2e−: 10 electronsA1

O + 2e− → O2−: 10 electronsA1

Quick check

1. Which noble gas has the same electron arrangement as K+?

2. Write the ion formed by nitrogen.

Answers: 1. argon 2. N3−

18 Precision in definitions

What examiners want.

Key definitions must be precise. Atomic number: the number of protons in the nucleus. Mass number: the number of protons plus neutrons. Isotopes: atoms of the same element with different numbers of neutrons. Relative atomic mass: weighted mean mass of an atom compared with 1/12 of a carbon-12 atom.

WORKED EXAMPLE 18 Correcting a definition (standard)

A student writes 'relative atomic mass is the mass of an atom'. Improve the definition.

Solution

Weighted mean mass of the atoms of an element (isotopes and abundances)R1

relative to 1/12 of the mass of a carbon-12 atomR1

Common mistake

Defining isotopes as 'elements with different numbers of neutrons'. Isotopes are atoms (or nuclides) of the same element; the word 'element' alone loses the mark.

19 Calculations with masses of particles

Real masses.

The mass of a proton or neutron is about 1.67 × 10−24 g and of an electron 9.11 × 10−28 g. The atomic mass unit (1 u) is 1/12 of the mass of a carbon-12 atom = 1.66 × 10−24 g.

WORKED EXAMPLE 19 Mass of an atom (standard)

Estimate the mass of one atom of 56Fe in grams.

Solution

56 × 1.66 × 10−24M1

= 9.30 × 10−23 gA1

WORKED EXAMPLE 20 Electron fraction (standard)

What percentage of the mass of a 12C atom is due to its electrons?

Solution

6 × 9.11 × 10−28/(12 × 1.66 × 10−24)M1

= 0.027%A1

Quick check

1. Estimate the number of atoms in 1.0 g of hydrogen-1 atoms.

Answers: 1. 1/(1.67 × 10−24) = 6.0 × 1023

20 Mixed review: worked problems

Bringing the ideas together.

WORKED EXAMPLE 21 Mixed problem (exam level)

Element X has two isotopes, AX (60%) and A+2X (40%), and Ar = 64.8. (a) Find A. (b) Identify X if it has 29 protons. (c) State the number of neutrons in the heavier isotope.

Solution

(a) 0.6A + 0.4(A + 2) = 64.8: A = 64M1 A1

(b) copper (although real copper is 63Cu/65Cu)A1

Quick check

1. Iridium has isotopes 191Ir and 193Ir and Ar = 192.22. Find the abundance of 191Ir.

Answers: 1. 39%

21 Additional exam-style practice

Paper 1 and Paper 2.

WORKED EXAMPLE 22 Paper 1 question (basic)

Which species has more electrons than neutrons? A. 21H B. 32He+ C. 168O2− D. 199F−

Solution

C: 10 electrons, 8 neutronsA1

WORKED EXAMPLE 23 Paper 2 question (exam level)

Neon has three isotopes: 20Ne (90.48%), 21Ne (0.27%), 22Ne (9.25%). (a) Calculate Ar to two decimal places. (b) Sketch its mass spectrum. (c) Explain why all three isotopes have the same chemical properties.

Solution

(a) 20.19M1 A1

(b) three peaks at 20, 21, 22 with heights in proportionA1

Quick check

1. State the number of neutrons in 235U.

Answers: 1. 143

22 Common misconceptions and links

Students often think…	Actually…
isotopes have different chemical properties	same electron arrangement, so same chemistry
Ar has units of grams	Ar has no units
ions have different numbers of protons	only electrons change

How S1.2 connects to the rest of the course

Topic	Connection
S1.3 Electron configurations	arrangement of electrons
S1.4 The mole	molar mass from Ar
S3.2 Functional groups	mass spectra of compounds
S3.1 Periodic table	atomic number and periodicity

23 Exam-style question with markscheme

Try this under exam conditions (about 15 minutes) before reading the markscheme.

Figure 3. Mass spectrum of magnesium.

(a) State the number of neutrons in the most abundant isotope. [1]

(b) Calculate the relative atomic mass of magnesium to two decimal places. [2]

(d) Deduce the numbers of protons and electrons in 25Mg2+. [2]

Markscheme

Part	Answer	Marks
(a)	12	A1
(b)	(24 × 79 + 25 × 10 + 26 × 11) ÷ 100	M1
	= 24.32	A1
(c)	same number of electrons	R1
	same electron configuration determines chemistry	R1
(d)	12 protons	A1
	10 electrons	A1

Examiner's comments. In (b), answers given to one decimal place or with units lost a mark.

24 Check your understanding

Section A: multiple choice

1. The atomic number is the number of

A. protons B. neutrons C. nucleons D. electrons in an ion

2. 4019K has how many neutrons?

A. 19 B. 40 C. 21 D. 59

3. An anion is formed by

A. losing electrons B. gaining electrons C. gaining protons D. losing neutrons

4. Isotopes differ in

A. number of protons B. chemical properties C. atomic number D. number of neutrons

5. In a mass spectrometer ions are separated by

A. colour B. mass to charge ratio C. boiling point D. solubility

6. Relative atomic mass is compared with

A. 1/12 of carbon-12 B. hydrogen-1 C. oxygen-16 D. 1 gram

Section B: structured

7. Describe the nuclear model of the atom. [3]

8. Copper is 69.2% 63Cu and 30.8% 65Cu. Calculate Ar. [2]

9. Outline the stages of mass spectrometry. [4]

Answers

1. A 2. C 3. B 4. D 5. B 6. A

7. Small dense positive nucleus of protons and neutrons; electrons around; mostly empty space.

8. (63 × 69.2 + 65 × 30.8) ÷ 100 = 63.62

9. Vaporization; ionization; acceleration; deflection by m/z; detection.

25 Summary and topic map

Idea	Remember
Nuclear model	dense positive nucleus; electrons around; Rutherford
Particles	proton +1, neutron 0, electron −1
Z and A	protons; protons + neutrons
Ions	electrons gained or lost
Isotopes	same Z, different neutrons
Ar	weighted average; no units
Mass spectrometry	m/z and abundance

Before the exam, make sure you…

•deduce particles from nuclear symbols
•calculate Ar to two decimal places
•explain isotopes' chemistry by electrons
•read peaks from spectra

The Nuclear Atom on one page

Cover this page with a sheet of paper and try to rebuild it from memory. Whatever you can't recall is what to revise next.

Figure 4. Topic map for The Nuclear Atom.

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