Atomic structure and the periodic table Revision Notes
Subject: Chemistry | Level: GCSE | Exam Board: AQA
Master the fundamental building blocks of all matter. This topic covers the structure of atoms, the history of atomic models, and how the periodic table predicts chemical properties—essential knowledge that underpins the entire GCSE Chemistry specification.
Revision Notes & Key Concepts
Revision Podcast Transcript
GCSE Chemistry Podcast: Atomic Structure and the Periodic Table Episode Script — Approximately 10 minutes [INTRO — 1 minute] Hello and welcome to your GCSE Chemistry revision podcast. I'm your tutor, and today we're diving into one of the most fundamental topics in the entire chemistry specification: Atomic Structure and the Periodic Table. This is topic 4.1, and it underpins almost everything else you'll study in chemistry — from bonding and reactions to the properties of materials. So if you nail this topic, you're building the strongest possible foundation for your exam success. Whether you're revising for the first time or doing a final check before your exam, this episode will walk you through every key concept, flag the mistakes that cost students marks every year, and finish with a quick-fire quiz so you can test yourself. Grab a pen and paper — you're going to want to take notes. Let's go. [CORE CONCEPTS — 5 minutes] Let's start right at the beginning: what actually is an atom? An atom is the smallest part of an element that can exist. Every single substance you've ever touched, breathed, or tasted is made of atoms. And every atom is made of three types of subatomic particle: protons, neutrons, and electrons. Here's the key table you absolutely must memorise. Protons have a relative mass of 1 and a relative charge of plus 1. Neutrons have a relative mass of 1 and a relative charge of zero — they're neutral, which is actually where the name comes from. And electrons have a relative mass of approximately 1 over 1840 — so tiny we treat it as zero — and a relative charge of minus 1. Now, where are these particles? Protons and neutrons are found in the nucleus at the centre of the atom. The nucleus is tiny but incredibly dense — it contains virtually all the mass of the atom. Electrons orbit the nucleus in shells, also called energy levels. Two numbers you must be able to use: the atomic number and the mass number. The atomic number — sometimes called the proton number — tells you how many protons are in the nucleus. This is the number that defines which element you're looking at. Change the number of protons and you've got a completely different element. The mass number tells you the total number of protons plus neutrons. So here's the calculation you'll be asked to do in the exam: to find the number of neutrons, you subtract the atomic number from the mass number. Mass number minus atomic number equals number of neutrons. For example, sodium has an atomic number of 11 and a mass number of 23. So the number of neutrons is 23 minus 11, which equals 12. Simple. And in a neutral atom — one with no overall charge — the number of electrons equals the number of protons. So sodium has 11 protons and 11 electrons. Now let's talk about isotopes. This is a definition that comes up in the exam regularly. Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because they have the same number of protons, they're the same element. But because they have different numbers of neutrons, they have different mass numbers. A brilliant example is carbon. Carbon-12 has 6 protons and 6 neutrons. Carbon-14 has 6 protons but 8 neutrons. Both are carbon — same atomic number of 6 — but different mass numbers. Carbon-14 is actually radioactive and is used in carbon dating. That's a lovely link to the radioactivity topic. Moving on to electronic structure. Electrons don't just float around randomly — they occupy specific energy levels, or shells. The first shell, closest to the nucleus, can hold a maximum of 2 electrons. The second shell can hold up to 8. The third shell can also hold up to 8 for the first 20 elements — which is all you need to know at GCSE. We write electronic configurations using numbers separated by commas. Sodium, with 11 electrons, has the configuration 2, 8, 1. That means 2 electrons in the first shell, 8 in the second, and 1 in the third. You can also draw these as Bohr diagrams — concentric circles with dots representing electrons. Here's the crucial link to the periodic table: the group number of an element tells you how many electrons are in its outer shell. So Group 1 elements have 1 outer electron. Group 7 elements have 7 outer electrons. Group 0 elements have a full outer shell — 8 electrons, or 2 for helium. This is why Group 0 elements are so unreactive: they already have a stable electron arrangement. Now let's talk about the periodic table itself. Elements are arranged in order of increasing atomic number — that's the number of protons. Each row is called a period, and each column is called a group. Elements in the same group have the same number of outer electrons, which is why they have similar chemical properties. Group 1 — the alkali metals — includes lithium, sodium, and potassium. They all have 1 outer electron, so they react by losing that electron to form a positive ion with a charge of plus 1. Reactivity increases as you go down Group 1. Why? Because as you go down, the outer electron is further from the nucleus, there are more inner shells shielding the nucleus's attraction, so the outer electron is lost more easily. Caesium, at the bottom of Group 1, is far more reactive than lithium at the top. Group 7 — the halogens — includes fluorine, chlorine, bromine, and iodine. They all have 7 outer electrons, so they react by gaining 1 electron to form a negative ion with a charge of minus 1. But here's the key difference: reactivity decreases as you go down Group 7. Fluorine is the most reactive halogen. As you go down, the outer shell is further from the nucleus and more shielded, so it's harder to attract an extra electron. This is the opposite trend to Group 1 — and mixing them up is one of the most common exam mistakes. The transition metals sit in the middle block of the periodic table. Compared to Group 1 metals, they are much harder, have much higher melting points, are much less reactive, and they form coloured compounds. They can also have variable oxidation states — for example, iron can be iron two plus or iron three plus. The term you should use when explaining their properties is delocalised electrons. Finally, let's briefly cover the development of the atomic model. Science doesn't stand still — our understanding of the atom has changed over time as new evidence emerged. Thomson proposed the plum pudding model in 1897: a positive sphere with electrons embedded in it, like plums in a pudding. Then Rutherford's gold foil experiment in 1911 showed that most of an atom is empty space with a tiny, dense, positive nucleus. This led to the nuclear model. Later, Bohr refined this by showing that electrons occupy fixed energy levels or shells. This is the model we use at GCSE. [EXAM TIPS AND COMMON MISTAKES — 2 minutes] Right, let's talk exam technique. Here are the mistakes that cost students marks every single year. Mistake number one: confusing atomic number with mass number. The atomic number is always the smaller number — it's the number of protons. The mass number is always the larger number — it's protons plus neutrons. In the exam, you'll be given a symbol with two numbers. The bottom number is always the atomic number. The top number is always the mass number. Mistake number two: getting the neutron calculation wrong. It's mass number MINUS atomic number. Not the other way around. Write it out as a formula: neutrons equals mass number minus atomic number. Mistake number three: mixing up Group 1 and Group 7 reactivity trends. Remember this: Group 1 reactivity INCREASES going down. Group 7 reactivity DECREASES going down. Group 1 metals are getting more explosive as you go down — caesium literally explodes in water. Group 7 halogens are getting less aggressive — iodine is far less reactive than fluorine. Mistake number four: drawing electronic configurations incorrectly. The first shell holds maximum 2. The second and third shells hold maximum 8. Never put more than 2 in the first shell. Never put more than 8 in the second shell. Mistake number five: not linking group number to outer electrons. If a question asks you to explain why elements in the same group have similar properties, you must say it's because they have the same number of outer shell electrons. That's the key phrase that earns the mark. For command words: if the question says 'State', give a brief factual answer — one or two words or a short phrase. If it says 'Explain', you must give a reason — use the word 'because' to link cause and effect. If it says 'Calculate', show your working, write the formula, substitute the values, and include units if required. Examiners are instructed to award marks for correct working even if the final answer is wrong. [QUICK-FIRE RECALL QUIZ — 1 minute] Time for your quick-fire quiz. I'll pause briefly after each question — try to answer before I give you the answer. Question 1: What is the relative charge of a proton? ... Plus 1. Question 2: An element has an atomic number of 17 and a mass number of 35. How many neutrons does it have? ... 35 minus 17 equals 18 neutrons. That's chlorine, by the way. Question 3: What is the electronic configuration of potassium, which has 19 electrons? ... 2, 8, 8, 1. Question 4: Reactivity in Group 1 — does it increase or decrease going down the group? ... It increases. Question 5: What is an isotope? ... Atoms of the same element with the same number of protons but different numbers of neutrons. How did you do? If you got all five, brilliant — you're in great shape. If you missed any, go back and review that section of your notes. [SUMMARY AND SIGN-OFF — 1 minute] Let's wrap up with the key points to take away from today's episode. One: atoms contain protons, neutrons, and electrons. Protons and neutrons are in the nucleus. Electrons are in shells around the nucleus. Two: atomic number equals number of protons. In a neutral atom, number of electrons equals number of protons. Number of neutrons equals mass number minus atomic number. Three: isotopes are atoms of the same element with different numbers of neutrons. Four: electronic configurations fill shells in order — maximum 2 in shell 1, maximum 8 in shells 2 and 3. Five: the group number tells you the number of outer electrons. This determines chemical properties. Six: Group 1 reactivity increases down the group. Group 7 reactivity decreases down the group. Group 0 elements are unreactive because they have a full outer shell. Seven: transition metals are harder, less reactive, have higher melting points, form coloured compounds, and have variable oxidation states compared to Group 1 metals. That's everything you need for Atomic Structure and the Periodic Table. You've got this. Keep revising, keep testing yourself, and I'll see you in the next episode. Good luck!
Key Terms & Definitions
- Atom
- The smallest part of an element that can exist.
- Isotope
- Atoms of the same element with the same number of protons but a different number of neutrons.
- Atomic Number
- The number of protons in the nucleus of an atom.
- Mass Number
- The total number of protons and neutrons in the nucleus of an atom.
- Ion
- An atom or group of atoms with a positive or negative charge, formed by the loss or gain of electrons.
- Electronic Configuration
- The arrangement of electrons in shells or energy levels around the nucleus of an atom.
Worked Examples
Worked Example
Question: An atom of potassium has an atomic number of 19 and a mass number of 39. Calculate the number of protons, neutrons, and electrons in this atom. (3 marks)
Solution: Step 1: The number of protons is equal to the atomic number. Protons = 19. Step 2: In a neutral atom, the number of electrons equals the number of protons. Electrons = 19. Step 3: The number of neutrons is the mass number minus the atomic number (39 - 19). Neutrons = 20. Final answer: 19 protons, 20 neutrons, 19 electrons.
Worked Example
Question: Explain why the reactivity of the halogens (Group 7) decreases as you go down the group. (4 marks)
Solution: Step 1: As you go down Group 7, the atoms gain more electron shells, so the outer shell is further from the nucleus. Step 2: This means there is increased shielding from inner electron shells. Step 3: Therefore, the electrostatic attraction between the positive nucleus and the incoming negative electron is weaker. Step 4: As a result, it is harder for the atom to gain the one electron needed to complete its outer shell, making it less reactive.
Worked Example
Question: Describe the differences between the plum pudding model of the atom and the nuclear model of the atom. (4 marks)
Solution: Step 1: In the plum pudding model, the atom is a solid sphere of positive charge, whereas in the nuclear model, the positive charge is concentrated in a tiny central nucleus. Step 2: In the plum pudding model, electrons are embedded randomly throughout the positive sphere, whereas in the nuclear model, electrons orbit the nucleus. Step 3: The nuclear model consists mostly of empty space, which is not a feature of the plum pudding model. Step 4: The plum pudding model does not have a distinct nucleus, whereas the nuclear model does.
Practice Questions
Question: Chlorine has two common isotopes: Chlorine-35 and Chlorine-37. Explain why these two isotopes have the same chemical properties. (2 marks)
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Question: Compare the chemical and physical properties of transition elements with Group 1 elements. (6 marks)
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Question: Describe how the alpha particle scattering experiment led to the nuclear model of the atom. (4 marks)
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Question: An element has the electronic structure 2,8,3. Identify its group and period in the periodic table. (2 marks)
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Question: Explain why noble gases (Group 0) are unreactive. (2 marks)
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