Welcome to Atomic Structure!
Everything in the universe — the screen you are reading this on, the air you breathe, and your own body — is made of tiny building blocks called atoms. Understanding how atoms are put together is the foundation for all of GCSE Chemistry. Don't worry if science has felt overwhelming before; we will break down every idea into clear, bite-sized steps so you feel fully confident for your exam!
---1. How the Model of the Atom Developed
Our idea of what an atom looks like did not appear overnight. Over the last two centuries, scientists conducted key experiments that changed our understanding step-by-step.
Key Milestones in Atomic History
1. John Dalton:
Dalton pictured atoms as solid, indivisible spheres (like tiny, solid billiard balls) that could not be split into anything smaller.
2. J.J. Thomson (1897) — The Plum Pudding Model:
Thomson discovered the first subatomic particle: the electron. Because electrons have a negative charge and atoms are neutral overall, he suggested the Plum Pudding Model. In this model, the atom is a sphere of positive charge with tiny, negatively charged electrons dotted throughout it (like plums in a pudding or chocolate chips in a muffin).
3. Ernest Rutherford — Alpha Particle Scattering Experiment:
Rutherford fired positively charged alpha particles at a very thin sheet of gold foil.
• Observation: Most alpha particles passed straight through the foil without changing direction.
• Conclusion: The atom is mostly empty space.
• Observation: A small number of alpha particles were deflected or bounced straight back.
• Conclusion: The positive charge and most of the atom's mass are concentrated in a tiny, dense center called the nucleus.
4. Niels Bohr:
Bohr refined Rutherford's model by calculating that electrons do not just float anywhere; they orbit the nucleus in fixed orbits or electron shells (energy levels) at set distances from the nucleus.
5. James Chadwick (1932):
Chadwick discovered the neutron. This particle had been difficult to detect because it carries no electrical charge. His discovery explained the "missing mass" in the nucleus.
Quick Summary of the Timeline:
Dalton (solid sphere) → Thomson (plum pudding / discovered electrons) → Rutherford (nuclear model / mostly empty space) → Bohr (electron shells) → Chadwick (discovered neutrons).
2. Subatomic Particles and Atomic Scale
An atom is made of three subatomic particles: protons, neutrons, and electrons.
Properties of Subatomic Particles
Proton (\(p\)):
• Relative Mass: \(1\)
• Relative Charge: \(+1\)
• Location: In the nucleus
Neutron (\(n\)):
• Relative Mass: \(1\)
• Relative Charge: \(0\) (neutral)
• Location: In the nucleus
Electron (\(e^-\)):
• Relative Mass: \(\frac{1}{1840}\) (negligible/negligibly small)
• Relative Charge: \(-1\)
• Location: In electron shells (orbits) surrounding the nucleus
Why do atoms have no overall charge?
Every neutral atom contains an equal number of positive protons and negative electrons. The positive charges and negative charges cancel each other out completely, leaving an overall charge of zero.
The Scale of Atoms and Nuclei
Atoms are incredibly tiny, and their nuclei are even smaller:
• Radius of an atom: Approximately \(0.1\text{ nm}\) (which equals \(1 \times 10^{-10}\text{ m}\)).
• Radius of a nucleus: Less than \(\frac{1}{10\,000}\) of the radius of the atom (less than \(1 \times 10^{-14}\text{ m}\)).
Real-World Analogy: If an atom were the size of a giant football stadium, the nucleus would be the size of a tiny marble sitting on the center spot, and the rest of the stadium would be empty space where the electrons orbit!
Memory Trick:
• Proton = Positive (\(+1\))
• Neutron = Neutral (\(0\))
• Electron = Negative (\(-1\))
3. Atomic Number, Mass Number, and Notation
Every element in the Periodic Table is identified by two key numbers.
Definitions
• Atomic Number (\(Z\)): The number of protons in the nucleus of an atom. This number identifies the element.
• Mass Number (\(A\)): The total number of protons and neutrons in the nucleus of an atom.
Standard Representation
We write elements using the symbol notation \(^{A}_{Z}\text{X}\):
• Top number (\(A\)) = Mass Number (larger number)
• Bottom number (\(Z\)) = Atomic Number (smaller number)
How to Calculate Subatomic Particles
When given any neutral atom \(^{A}_{Z}\text{X}\):
1. \(\text{Number of protons} = \text{Atomic Number } (Z)\)
2. \(\text{Number of electrons} = \text{Number of protons } (Z)\)
3. \(\text{Number of neutrons} = \text{Mass Number } (A) - \text{Atomic Number } (Z)\)
Worked Example:
Let's find the subatomic particles in Sodium: \(^{23}_{11}\text{Na}\)
• Protons = \(11\)
• Electrons = \(11\)
• Neutrons = \(23 - 11 = 12\)
4. Electronic Configuration (Arranging Electrons)
Electrons occupy specific energy levels or electron shells around the nucleus. They always fill the lowest available energy shell (closest to the nucleus) first.
Shell Rules for the First 20 Elements:
• 1st Shell: Holds up to 2 electrons
• 2nd Shell: Holds up to 8 electrons
• 3rd Shell: Holds up to 8 electrons
• 4th Shell: Fills up to 2 electrons (for Calcium, atomic number \(20\))
Writing and Drawing Configurations
In CCEA GCSE, electronic configurations are written as numbers separated by commas:
• Carbon (\(6\) electrons): \(2,4\)
• Sodium (\(11\) electrons): \(2,8,1\)
• Chlorine (\(17\) electrons): \(2,8,7\)
• Calcium (\(20\) electrons): \(2,8,8,2\)
When drawing Bohr circle diagrams, represent electrons on each ring using dots (•) or crosses (\(\times\)). Pair up electrons once there are more than 4 in a shell to make counting easy.
Linking Electron Structure to the Periodic Table
• Group Number = The number of electrons in the outer shell (e.g., Sodium is \(2,8,1\); it has 1 outer electron, so it is in Group 1).
• Period Number = The number of occupied electron shells (e.g., Sodium has 3 shells containing electrons, so it is in Period 3).
5. Isotopes and Relative Atomic Mass (\(A_r\))
What is an Isotope?
Isotopes are atoms of the same element with the same number of protons (same atomic number) but a different number of neutrons (different mass number).
Example: Carbon Isotopes
• Carbon-12 (\(^{12}_{6}\text{C}\)): \(6\) protons, \(6\) electrons, \(6\) neutrons.
• Carbon-14 (\(^{14}_{6}\text{C}\)): \(6\) protons, \(6\) electrons, \(8\) neutrons.
Why do isotopes have identical chemical properties?
Chemical reactions depend on the arrangement of electrons — especially the outer shell electrons. Because isotopes of an element have the exact same electronic configuration, they react in exactly the same way chemically.
Calculating Relative Atomic Mass (\(A_r\))
The Relative Atomic Mass (\(A_r\)) is the weighted average mass of the atoms of an element compared to \(\frac{1}{12}\)th the mass of a carbon-12 atom. It accounts for all natural isotopes and their percentage abundances.
Formula:
\(A_r = \frac{\sum (\text{Isotope Mass} \times \% \text{ Abundance})}{100}\)
Step-by-Step Worked Example:
A sample of chlorine contains \(75\%\) \(^{35}\text{Cl}\) and \(25\%\) \(^{37}\text{Cl}\). Calculate the relative atomic mass of chlorine.
Step 1: Multiply each mass by its percentage abundance:
\((35 \times 75) = 2625\)
\((37 \times 25) = 925\)
Step 2: Add these values together:
\(2625 + 925 = 3550\)
Step 3: Divide by 100:
\(A_r = \frac{3550}{100} = 35.5\)
Note: Always give your final answer to one decimal place if specified in an exam question!
---6. Ions
An ion is an atom or group of atoms that has gained or lost electrons, resulting in an overall electrical charge.
Why do atoms form ions?
Atoms gain or lose electrons to achieve a full outer shell of electrons, which gives them the stable electron structure of a noble gas.
Types of Ions
1. Positive Ions (Cations):
• Formed when metal atoms lose electrons.
• Example: A Sodium atom (\(2,8,1\)) loses \(1\) electron to form a sodium ion (\(\text{Na}^+\)) with configuration \(2,8\).
\(\text{Na} \rightarrow \text{Na}^+ + e^-\)
2. Negative Ions (Anions):
• Formed when non-metal atoms gain electrons.
• The name of a simple non-metal ion changes its ending to -ide.
• Example: A Chlorine atom (\(2,8,7\)) gains \(1\) electron to form a chloride ion (\(\text{Cl}^-\)) with configuration \(2,8,8\).
\(\text{Cl} + e^- \rightarrow \text{Cl}^-\)
Counting Subatomic Particles in Ions
Remember: The number of protons and neutrons in the nucleus NEVER changes when an ion forms! Only electrons are transferred.
• Positive charge (\(+n\)): \(\text{Electrons} = \text{Protons} - n\)
• Negative charge (\(-n\)): \(\text{Electrons} = \text{Protons} + n\)
Examples:
• \(\text{Mg}^{2+}\) (Magnesium: atomic number \(12\), mass number \(24\)): Protons = \(12\), Neutrons = \(12\), Electrons = \(12 - 2 = 10\).
• \(\text{O}^{2-}\) (Oxide: atomic number \(8\), mass number \(16\)): Protons = \(8\), Neutrons = \(8\), Electrons = \(8 + 2 = 10\).
7. Common Exam Pitfalls to Avoid
• Mass of an electron: Never write that an electron has a mass of "0". Write \(\frac{1}{1840}\), negligible, or negligibly small.
• Isotope definition: Always specify "atoms of the same element with the same number of protons but different numbers of neutrons". Don't just say "elements with different neutrons".
• Biology vs Chemistry Nucleus: An atomic nucleus contains protons and neutrons. Never confuse it with a biological cell nucleus (which contains DNA)!
• Ion charges: Remember that gaining negative electrons makes an ion negative, while losing electrons leaves behind a positive charge.
• Name vs Symbol: Pay close attention to command words. If the question asks to "Name the element", write "Sodium", not "Na".
Quick Revision Checklist
• Can you describe how the atomic model changed from Dalton to Chadwick?
• Do you know the relative mass and charge of protons, neutrons, and electrons?
• Can you calculate protons, neutrons, and electrons for any atom or ion?
• Can you write the electronic configuration for any of the first 20 elements?
• Can you calculate \(A_r\) using isotopic percentages?