Welcome to Atomic Structure!

Welcome to your first step into AQA A-Level Chemistry! This chapter is the foundation of everything else you will learn. Atoms are like the "Lego bricks" of the universe. If you understand how they are built, you can understand how they react, why they bond, and why the Periodic Table is shaped the way it is.
Don't worry if this seems like a lot to take in at first—we’ll break it down piece by piece!

1. The Fundamental Particles

At the heart of every atom are three sub-atomic particles. Even though atoms are incredibly small, these particles determine the identity and behavior of every element.

What’s inside the atom?

  • The Nucleus: This is the tiny, dense center of the atom containing protons and neutrons. Almost all the atom's mass is here!
  • The Electrons: These are tiny particles that "whiz" around the nucleus in specific areas called orbitals.

Mass and Charge Table

To keep things simple, we use "relative" mass and charge rather than the tiny actual numbers.

Proton: Relative Mass = 1 | Relative Charge = +1
Neutron: Relative Mass = 1 | Relative Charge = 0 (Neutral)
Electron: Relative Mass = 1/1840 (basically zero!) | Relative Charge = -1

Did you know? If an atom were expanded to the size of a football stadium, the nucleus would be the size of a small fly in the center, and the electrons would be like gnats buzzing around the very top seats! Atoms are mostly empty space.

Key Takeaway:

The nucleus (protons + neutrons) gives the atom its mass, while the electrons determine its chemical reactions.

2. Mass Number, Atomic Number, and Isotopes

Every element in the Periodic Table is defined by its numbers. You'll see two numbers next to an element's symbol, like \( ^{12}_{6}C \).

The Two Big Numbers:

  • Atomic (Proton) Number (Z): The number of protons. This is the "ID card" of the element. If you change the proton number, you change the element!
  • Mass Number (A): The total number of protons + neutrons.

Quick Trick to Remember:
A is the All-together number (Protons + Neutrons).
Z is the Zero-in-on-identity number (Protons only).

What are Isotopes?

Isotopes are atoms of the same element with the same number of protons but a different number of neutrons.
Because they have the same number of electrons, isotopes react identically in chemical reactions. They just have slightly different masses!

Common Mistake to Avoid: Students often think isotopes have different chemical properties. They don't! Chemistry is all about electrons, and isotopes have the same number of those.

Key Takeaway:

Neutrons = Mass Number (A) - Atomic Number (Z). For a neutral atom, Electrons = Protons.

3. Time of Flight (TOF) Mass Spectrometry

How do chemists weigh something as tiny as an atom? They use a Mass Spectrometer. In the AQA syllabus, we focus on the Time of Flight (TOF) method.

The Five Stages of TOF:

Imagine a race where the "prize" is hitting a detector at the end of a tube.

  1. Ionisation: The sample must be turned into positive ions. There are two ways:
    • Electron Impact: High-energy electrons are fired at the sample, knocking an electron off: \( X(g) \rightarrow X^+(g) + e^- \). (Used for elements/small molecules).
    • Electrospray: The sample is dissolved and pushed through a needle at high voltage, gaining a proton: \( X(g) + H^+ \rightarrow XH^+(g) \). (Used for big molecules like proteins).
  2. Acceleration: The positive ions are attracted toward a negative plate. They are all given the same kinetic energy.
  3. Ion Drift: The ions enter a "flight tube" with no electric field. They just drift along.
  4. Detection: The ions hit a negatively charged detector plate. When they hit, they gain an electron, which creates a current. The size of the current tells us how many ions arrived.
  5. Data Analysis: The computer records the "Time of Flight."

The "Race" Analogy:

If you give a bowling ball and a tennis ball the same "push" (Kinetic Energy), which one reaches the wall first? The tennis ball!
In the spectrometer, lighter ions travel faster and reach the detector in less time than heavier ions.

Calculations:

You can calculate the Relative Atomic Mass (\(A_r\)) using the data from a mass spectrum:
\( A_r = \frac{\sum (isotopic\ mass \times abundance)}{total\ abundance} \)

Key Takeaway:

TOF separates ions based on their mass-to-charge ratio (\(m/z\)). Since the charge is usually +1, it effectively separates them by mass.

4. Electron Configuration

Electrons aren't just a messy cloud; they are highly organized into shells and sub-shells.

The Levels of Organization:

  • Shells: The main energy levels (1, 2, 3, 4).
  • Sub-shells: Divided into s, p, d, and f.
  • Orbitals: Regions of space where an electron is likely to be found. Each orbital can hold exactly two electrons with opposite spins.

Capacities to Remember:

  • s sub-shell: 1 orbital (holds 2 electrons)
  • p sub-shell: 3 orbitals (holds 6 electrons)
  • d sub-shell: 5 orbitals (holds 10 electrons)

The Filling Order:

Electrons fill the lowest energy levels first. The order is:
1s \(\rightarrow\) 2s \(\rightarrow\) 2p \(\rightarrow\) 3s \(\rightarrow\) 3p \(\rightarrow\) 4s \(\rightarrow\) 3d

Wait! Why 4s before 3d?
The 4s sub-shell actually has a slightly lower energy than the 3d sub-shell, so it fills up first. Think of it like a hotel where the 4th floor "budget room" is cheaper than the 3rd floor "luxury suite."

The Two "Rule Breakers" (Exceptions):

AQA expects you to know Chromium (Z=24) and Copper (Z=29). They prefer to have a half-full or completely full d-subshell because it's more stable.
Cr: \( [Ar] 4s^1 3d^5 \) (not \( 4s^2 3d^4 \))
Cu: \( [Ar] 4s^1 3d^{10} \) (not \( 4s^2 3d^9 \))

Key Takeaway:

When atoms turn into ions, they lose electrons from the 4s orbital before the 3d orbital. This is a very common exam trap!

5. Ionisation Energy

First Ionisation Energy 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: \( X(g) \rightarrow X^+(g) + e^- \)

What affects Ionisation Energy?

  1. Nuclear Charge: More protons = more attraction = higher energy needed.
  2. Distance: Further from nucleus = less attraction = lower energy needed.
  3. Shielding: More inner shells "block" the pull of the nucleus = lower energy needed.

Trends in the Periodic Table:

  • Down a Group: Decreases. Even though there are more protons, the outer electron is much further away and has more shielding.
  • Across a Period: Generally increases. The nuclear charge increases (more protons) while shielding stays roughly the same.

The Evidence for Sub-shells:

If you look at Period 3 (Na to Ar), there are two "dips" where the energy drops slightly. These dips are the "smoking gun" that proves sub-shells exist!

  • Dip at Magnesium to Aluminium: The electron removed from Al is in a 3p orbital, which is higher in energy and slightly shielded by the 3s, making it easier to remove.
  • Dip at Phosphorus to Sulfur: In Sulfur, the electron is removed from a p-orbital containing two electrons. These two electrons repel each other, making it easier to "kick one out."
Key Takeaway:

Successive ionisation energies (removing the 2nd, 3rd, 4th electron) show big jumps when you move into a new shell closer to the nucleus.

Quick Review Quiz:

  • Which particle determines the element's identity? (Proton)
  • Why do isotopes have the same chemical properties? (Same electron configuration)
  • In TOF, which hits the detector first: \( ^{24}Mg^+ \) or \( ^{25}Mg^+ \)? (\( ^{24}Mg^+ \) because it is lighter)
  • Which orbital fills after 3p? (4s)

Well done! You've covered the essentials of AQA Atomic Structure. Take a break, grab a drink, and come back to practice some mass spectrometry calculations!