Welcome to Atomic Structure!
Welcome to the world of the ultra-small! In this chapter, we are going exploring the building blocks of everything you see around you. Understanding the atom is like learning the alphabet of Chemistry—once you know how atoms are put together, the rest of the subject starts to make much more sense. Atomic structure is part of your AS Level Physical Chemistry journey, and it’s all about where particles live inside an atom and how they behave.
Don't worry if some of this feels abstract at first. We will use plenty of analogies and step-by-step guides to make sure you feel confident!
1.1 Particles in the Atom
If you were to enlarge an atom to the size of a massive football stadium, the nucleus (the center) would be no bigger than a small marble in the middle of the pitch. The rest of the stadium? That’s mostly empty space!
The Subatomic "Team"
There are three main particles you need to know. Think of them by their mass and their "attitude" (charge):
- Protons: Heavy and positive. They live in the nucleus.
- Neutrons: Heavy and neutral (no charge). They also live in the nucleus and act like "glue."
- Electrons: Extremely light and negative. They zip around the nucleus in shells.
Relative Mass and Charge
In Chemistry, we use "relative" values because the actual grams and coulombs are too tiny to work with easily.
Proton: Mass = 1 | Charge = +1
Neutron: Mass = 1 | Charge = 0
Electron: Mass = 1/1840 (almost zero) | Charge = -1
The Atom's ID Card: Atomic and Mass Numbers
We use the notation \( ^x_yA \) to identify atoms:
- Atomic Number (\( y \)): Also called the proton number. This defines the element. If you change this, you change the element!
- Mass Number (\( x \)): Also called the nucleon number. This is the total number of protons + neutrons.
Quick Review: How to calculate particles
1. Protons = Atomic Number
2. Electrons = Protons (in a neutral atom)
3. Neutrons = Mass Number \( - \) Atomic Number
Beams in an Electric Field
If you fire these particles between a positive and a negative plate:
1. Protons curve toward the negative plate.
2. Electrons curve sharply toward the positive plate (they curve more because they are much lighter).
3. Neutrons go straight through—they don't care about charges!
Key Takeaway: Atoms are mostly empty space. The mass is concentrated in the tiny, positive nucleus, while the negative electrons occupy the space around it.
1.2 Isotopes
Imagine two versions of the same smartphone: one has a standard battery, and the other has a heavy-duty battery. They look and function the same, but one is heavier. These are like isotopes.
Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons.
Properties of Isotopes
- Chemical Properties: These stay the SAME. Chemical reactions depend on electrons, and isotopes have the same number of electrons.
- Physical Properties: These are DIFFERENT. Because they have different numbers of neutrons, isotopes have different masses and densities.
Did you know? Most Hydrogen atoms have no neutrons, but a rare isotope called Deuterium has one. It makes "heavy water," which actually sinks in normal water!
1.3 Electrons: Energy Levels and Orbitals
Electrons don't just fly around randomly; they live in a very organized "apartment building" system.
The Hierarchy
1. Shells: The main levels (Principal Quantum Number, \( n = 1, 2, 3... \)).
2. Sub-shells: Named s, p, d, and f.
3. Orbitals: The actual "rooms" where electrons live. Each orbital can hold at most 2 electrons.
Shapes to Know
- s orbitals: Spherical shape.
- p orbitals: Dumbbell shape (there are three types: \( p_x, p_y, p_z \)).
The Filling Order (The Aufbau Principle)
Electrons fill the lowest energy levels first. The order is:
\( 1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \)
Wait! Notice that the 4s sub-shell fills before the 3d sub-shell because it is slightly lower in energy. However, when writing the full configuration for ions, remember that 4s electrons are also the first to leave!
Electronic Configuration Example: Iron (Fe, Atomic Number 26)
Full: \( 1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2 \)
Shorthand: \( [Ar] 3d^6 4s^2 \)
Common Mistake: When drawing "electrons in boxes," remember Hund's Rule: electrons prefer to occupy their own orbital singly before pairing up. It's like people sitting on a bus—they usually take an empty double seat before sitting next to a stranger!
Free Radicals: These are species with one or more unpaired electrons. They are very reactive because that "lonely" electron wants a partner!
Key Takeaway: Electrons occupy specific orbitals. The 4s sub-shell is lower in energy than 3d, which is why it fills first.
1.4 Ionisation Energy (IE)
Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions.
The Equation (First IE)
\( X(g) \rightarrow X^+(g) + e^- \)
Note: Always include the (g) state symbol. Examiners look for this!
Factors Affecting IE
Think of it as a "tug-of-war" between the nucleus and the electron:
- Nuclear Charge: More protons = stronger pull on electrons (Higher IE).
- Atomic Radius: Electron further away = weaker pull (Lower IE).
- Shielding: More inner shells "block" the pull of the nucleus (Lower IE).
- Spin-pair Repulsion: Two electrons in the same orbital push each other away, making it easier to remove one (Lower IE).
Trends in the Periodic Table
- Across a Period: IE generally increases. The nuclear charge increases and the radius decreases, while shielding stays similar.
- Down a Group: IE decreases. Even though there are more protons, the extra shells increase the distance and shielding significantly.
Successive Ionisation Energies
You can remove electrons one by one (1st, 2nd, 3rd...). If you see a huge jump in energy, it means you've started removing an electron from a new shell that is closer to the nucleus.
Example: If the 1st and 2nd IEs are low, but the 3rd is massive, the element has 2 electrons in its outer shell (it’s in Group 2).
Key Takeaway: Ionisation energy tells us how strongly an atom holds onto its electrons. Jumps in successive IE data help us identify the Group an element belongs to.
Quick Review Summary
1. Particles: Protons (+) and Neutrons (0) in the nucleus; Electrons (-) in shells.
2. Isotopes: Same protons, different neutrons. Same chemistry, different mass.
3. Orbitals: \( s \) is a sphere, \( p \) is a dumbbell. Fill 4s before 3d.
4. IE Trends: Increases across a period, decreases down a group. Watch for the "big jump" to find the Group number.