Welcome to Atomic Structure!
Welcome to one of the most fundamental chapters in AS Level Chemistry! Have you ever wondered why different elements behave in completely unique ways? Why does sodium explode in water while neon does absolutely nothing? The secret lies entirely in atomic structure: the tiny particles inside atoms and how electrons arrange themselves.
Don't worry if you find this topic a bit abstract at first. We will break down every single concept into small, bite-sized steps with clear analogies and examples so that you feel fully confident for your CCEA AS 1 exam.
1. Inside the Atom: Subatomic Particles
Atoms are the building blocks of all matter, but they are made of three key subatomic particles:
• Protons: Located in the nucleus. Relative mass = \(1\), Relative charge = \(+1\).
• Neutrons: Located in the nucleus. Relative mass = \(1\), Relative charge = \(0\) (neutral).
• Electrons: Orbiting the nucleus in shells/orbitals. Relative mass = \(\frac{1}{1840}\) (often written as negligible or \(0.00055\)), Relative charge = \(-1\).
Atomic Number and Mass Number
Every element in the Periodic Table is defined by two numbers:
• Atomic Number (\(Z\)): The number of protons in the nucleus of an atom. This defines which element it is! In a neutral atom, the number of protons equals the number of electrons.
• Mass Number (\(A\)): The total number of protons plus neutrons in the nucleus.
• To find the number of neutrons: \(\text{Number of neutrons} = A - Z\).
What Are Isotopes?
Definition: 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).
Did you know? Isotopes of the same element have identical chemical properties because they have the same electron configuration! However, they have slightly different physical properties (such as density, melting point, or rate of diffusion) because their masses are different.
Key Takeaway: Protons define the element, neutrons determine the isotope, and electrons govern chemical reactions!
2. Relative Masses and Mass Spectrometry
Important Definitions
Because atoms are incredibly small, chemists measure their masses relative to a standard:
• Relative Atomic Mass (\(A_r\)): The weighted mean mass of an atom of an element relative to \(\frac{1}{12}\)th of the mass of an atom of carbon-12.
• Relative Isotopic Mass: The mass of an atom of an isotope relative to \(\frac{1}{12}\)th of the mass of an atom of carbon-12.
How a Mass Spectrometer Works
A mass spectrometer is an instrument used to determine the relative isotopic masses and their percentage abundances in an element. Think of it like a gust of wind blowing tennis balls and bowling balls: lighter particles are deflected or travel faster than heavier ones!
The four fundamental stages are:
1. Ionisation: The sample is vaporised and converted into positive ions (usually by knocking off an electron).
2. Acceleration: The positive ions are accelerated by an electric field so that they all have the same kinetic energy.
3. Deflection / Drift: Ions are separated based on their mass-to-charge ratio (\(m/z\)). Lighter ions (or those with higher charge) are deflected more / travel faster.
4. Detection: Ions hit a detector, producing an electric current. The size of the current is directly proportional to the abundance of that specific ion.
Calculating Relative Atomic Mass (\(A_r\))
To calculate the \(A_r\), use the following formula:
\(A_r = \frac{\sum (\text{isotopic mass} \times \text{abundance})}{\text{total abundance}}\)
Worked Example:
A sample of chlorine contains \(75\%\) of \(^{35}\text{Cl}\) and \(25\%\) of \(^{37}\text{Cl}\).
\(A_r = \frac{(35 \times 75) + (37 \times 25)}{100} = \frac{2625 + 925}{100} = \frac{3550}{100} = 35.5\)
Common Mistake to Avoid: Always check if the abundances are given as percentages (which add to \(100\)) or as relative peak heights/ratios (which add to a different total). Always divide by the total abundance!
Key Takeaway: Mass spectrometry gives us the isotopic composition of elements, allowing us to find the weighted average mass (\(A_r\)).
3. Electron Configuration and Orbitals
At GCSE, you learned that electrons sit in shells (\(2, 8, 8\)). At AS Level, we zoom in closer! Shells are split into sub-shells, and sub-shells are made of orbitals.
What is an Orbital?
Definition: An atomic orbital is a region of space around the nucleus where there is a high probability (over \(95\%\)) of finding an electron. Each single orbital can hold a maximum of two electrons with opposite spins.
Sub-shells and Shapes
• \(s\)-subshell: Has \(1\) spherical orbital \(\rightarrow\) holds up to \(2\) electrons.
• \(p\)-subshell: Has \(3\) dumbbell-shaped orbitals (\(p_x, p_y, p_z\)) \(\rightarrow\) holds up to \(6\) electrons.
• \(d\)-subshell: Has \(5\) orbitals \(\rightarrow\) holds up to \(10\) electrons.
• \(f\)-subshell: Has \(7\) orbitals \(\rightarrow\) holds up to \(14\) electrons.
The Energy Level Order
Electrons fill the lowest available energy levels first (the Aufbau Principle):
\(1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p\)
Crucial Point: Notice that the \(4s\) sub-shell has lower energy than the \(3d\) sub-shell, so \(4s\) fills before \(3d\)!
Rules for Filling Orbitals
1. Aufbau Principle: Fill the lowest energy orbital first.
2. Pauli Exclusion Principle: An orbital holds at most two electrons, and they must have opposite spins (represented by \(\uparrow\downarrow\)).
3. Hund's Rule: Electrons occupy orbitals of the same energy singly before pairing up (just like people picking empty double seats on a bus before sitting next to a stranger!).
Two Special Exceptions: Chromium and Copper
Nature loves stability! A half-filled or fully-filled \(d\)-subshell provides extra stability:
• Chromium (\(Z = 24\)): \([\text{Ar}]\, 4s^1 3d^5\) (NOT \(4s^2 3d^4\))
• Copper (\(Z = 29\)): \([\text{Ar}]\, 4s^1 3d^{10}\) (NOT \(4s^2 3d^9\))
Writing Electron Configurations for Ions
• For positive ions (cations): Remove electrons.
• CRITICAL RULE for transition metals: When forming transition metal ions, electrons are lost from the \(4s\) sub-shell before the \(3d\) sub-shell! (First in, first out).
Example: \(\text{Fe} = 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\)
\(\text{Fe}^{2+} = 1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\) (the two \(4s\) electrons are removed!).
Key Takeaway: \(4s\) fills before \(3d\), but when making ions, \(4s\) empties before \(3d\)!
4. Ionisation Energy
Definition of First Ionisation Energy
Definition: The first ionisation energy (\(\text{IE}_1\)) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous \(1+\) ions.
Equation: \(X\text{(g)} \rightarrow X^+\text{(g)} + e^-\)
Important: Always include the state symbols \(\text{(g)}\) in ionisation equations!
Successive Ionisation Energies
• Second Ionisation Energy: \(X^+\text{(g)} \rightarrow X^{2+}\text{(g)} + e^-\)
• Third Ionisation Energy: \(X^{2+}\text{(g)} \rightarrow X^{3+}\text{(g)} + e^-\)
Successive ionisation energies always increase because you are removing a negative electron from an increasingly positive ion, which pulls the remaining electrons tighter.
Three Key Factors that Determine Ionisation Energy
Whenever an exam question asks you to explain differences in ionisation energy, structure your answer around these three factors:
1. Nuclear Charge: More protons = greater positive charge = stronger attraction for outer electrons = higher IE.
2. Atomic Radius (Distance): Greater distance from nucleus = weaker attraction = lower IE.
3. Shielding: More inner electron shells = repulsion of outer electrons from inner electrons = weaker attraction to the nucleus = lower IE.
Trends in First Ionisation Energy
1. Down a Group (Decreases)
• Atomic radius increases (more shells).
• Shielding increases significantly.
• Even though nuclear charge increases, the extra shielding and distance outweigh it.
• Therefore, outer electron is less strongly held \(\rightarrow\) First IE decreases down a group.
2. Across a Period (General Increase)
• Number of protons increases (higher nuclear charge).
• Electrons are added to the same main energy level, so shielding remains roughly constant.
• Atomic radius decreases because the stronger nuclear charge pulls the shell closer.
• Therefore, outer electron is held more tightly \(\rightarrow\) First IE generally increases across a period.
The Two Discontinuities Across Period 2 and Period 3
Examiners love testing the two small "dips" in the across-period trend. Let's look at Period 3 (Mg to Al and P to S):
Dip 1: Magnesium to Aluminium (Group 2 to Group 3)
• \(\text{Mg}: 1s^2 2s^2 2p^6 3s^2\)
• \(\text{Al}: 1s^2 2s^2 2p^6 3s^2 3p^1\)
• Explanation: The electron removed from \(\text{Al}\) is in a \(3p\) sub-shell, which is higher in energy and shielded by the \(3s\) electrons. It is therefore easier to remove than the \(3s\) electron in \(\text{Mg}\).
Dip 2: Phosphorus to Sulfur (Group 5 to Group 6)
• \(\text{P}: 1s^2 2s^2 2p^6 3s^2 3p_x^1 3p_y^1 3p_z^1\)
• \(\text{S}: 1s^2 2s^2 2p^6 3s^2 3p_x^2 3p_y^1 3p_z^1\)
• Explanation: In \(\text{S}\), one of the \(3p\) orbitals contains a pair of electrons. Spin-pair repulsion between the two electrons in the same orbital makes it easier to remove one of them compared to the singly-occupied orbitals in \(\text{P}\).
Using Successive Ionisation Energies to Find Group Numbers
A huge jump in successive ionisation energies indicates that an electron is being removed from an inner, complete shell closer to the nucleus.
Example: An element has the following successive IEs (in \(\text{kJ mol}^{-1}\)):
\(578, 1817, 2745, \mathbf{11578}, 14842\)
Notice the enormous jump between the 3rd and 4th ionisation energy! This means \(3\) electrons were in the outer shell before breaking into a new inner shell. Therefore, the element is in Group 3.
Key Takeaway: Nuclear charge, distance, and shielding explain all IE trends. A giant jump in successive IE values tells you the group number!
Quick Chapter Summary Checklist
Before moving on to the next chapter, check that you can:
• State the charge, mass, and location of protons, neutrons, and electrons.
• Define isotopes, relative atomic mass (\(A_r\)), and relative isotopic mass.
• Calculate \(A_r\) from mass spectrometry abundance data.
• Write full \(s, p, d\) electron configurations for atoms and ions up to \(Z=36\) (remembering \(\text{Cr}\) and \(\text{Cu}\)).
• Write the balanced equation for the 1st ionisation energy of any element with state symbols.
• Explain the trends in ionisation energy across a period and down a group, including the two anomalies.