Welcome to Atomic Structure and the Periodic Table!
Welcome to the very beginning of your A Level Chemistry journey! This chapter is part of Paper 1: Advanced Inorganic and Physical Chemistry. Think of this topic as the "Instruction Manual" for the universe. Before we can understand how chemicals react or why materials behave the way they do, we need to look at the tiny building blocks they are made of: atoms.
Don't worry if some of this feels abstract at first. We will break down the "invisible" world of subatomic particles into simple, logical steps. By the end of these notes, you'll be able to predict the behavior of elements just by looking at their position on the Periodic Table!
1. The Structure of the Atom
Atoms are the smallest parts of an element that can exist. Even though they are tiny, they are made of even smaller sub-atomic particles. You need to know three: protons, neutrons, and electrons.
Mass and Charge
In Chemistry, we use "relative" mass and charge because the actual numbers are too tiny to work with easily.
Proton: Relative Mass = 1 | Relative Charge = +1
Neutron: Relative Mass = 1 | Relative Charge = 0 (neutral)
Electron: Relative Mass = 1/1840 (almost zero!) | Relative Charge = -1
Atomic and Mass Numbers
Atomic (Proton) Number (Z): The number of protons in the nucleus. This defines the element. If you change the number of protons, you change the element!
Mass Number (A): The total number of protons + neutrons in the nucleus.
Analogy: Think of the Atomic Number as an element's ID card—it never changes for that element. The Mass Number is like the element's weight on a scale.
Calculating Particles
To find the number of each particle in an atom:
1. Protons = Atomic Number.
2. Electrons = Same as protons (in a neutral atom).
3. Neutrons = Mass Number \( - \) Atomic Number.
Common Mistake to Avoid: When dealing with ions, the number of protons stays the same. Only the number of electrons changes. A positive ion (cation) has lost electrons, and a negative ion (anion) has gained them.
Key Takeaway: Atoms are made of protons and neutrons in a tiny central nucleus, with electrons orbiting in shells. The atomic number tells you "who" the element is.
2. Isotopes and Relative Mass
Not all atoms of the same element are identical. Some have more "baggage" than others.
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 chemically in the exact same way. However, they have different physical properties (like density) because their masses are different.
The Carbon-12 Scale
Everything in atomic mass is compared to Carbon-12.
Relative Isotopic Mass: The mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12.
Relative Atomic Mass (Ar): The weighted mean mass of an atom of an element compared with 1/12th of the mass of an atom of carbon-12.
Quick Review: Calculating Ar
To calculate the Relative Atomic Mass from isotopic abundance:
\( Ar = \frac{\sum (isotopic \ mass \ \times \ percentage \ abundance)}{100} \)
Did you know? We use Carbon-12 as the standard because it is a common solid and easy to transport and measure accurately in a lab!
Mass Spectrometry
A Mass Spectrometer is a machine that tells us the mass and abundance of isotopes in a sample.
1. For atoms, the peak furthest to the right (highest m/z) tells you the relative isotopic mass.
2. For molecules, the peak with the highest m/z value is the molecular ion peak (M+), which gives the Relative Molecular Mass of the whole molecule.
The Chlorine Molecule (Cl2) Mystery: Chlorine has two main isotopes: \( ^{35}Cl \) and \( ^{37}Cl \). In a mass spectrum of \( Cl_2 \), you will see three peaks at m/z 70, 72, and 74. This is because the molecules can be: \( ^{35}Cl-^{35}Cl \), \( ^{35}Cl-^{37}Cl \), or \( ^{37}Cl-^{37}Cl \).
Key Takeaway: Isotopes are versions of an element with different neutron counts. Ar is the "average" mass of all those versions combined.
3. Ionisation Energy
First Ionisation Energy: The energy required to remove one mole of electrons from one mole of gaseous atoms to provide one mole of gaseous 1+ ions.
Equation: \( X(g) \rightarrow X^+(g) + e^- \)
Factors Affecting Ionisation Energy (The "Big Three")
If you find this tricky, always go back to these three points:
1. Nuclear Charge: More protons in the nucleus = more "pull" on electrons = higher IE.
2. Atomic Radius: Electrons further from the nucleus are held less tightly = lower IE.
3. Shielding: Inner shells of electrons block the "pull" of the nucleus on outer electrons = lower IE.
Trends in First Ionisation Energy
Across a Period (Left to Right): IE increases. The nuclear charge increases (more protons) while shielding stays roughly the same. The "pull" gets stronger.
Down a Group: IE decreases. Even though there are more protons, the atomic radius increases and there is more shielding. The outer electron is much easier to steal!
Successive Ionisation Energies: You can keep removing electrons (2nd, 3rd, etc.). A large jump in IE tells you that an electron has been removed from a shell closer to the nucleus. This helps us identify which Group an element belongs to.
Key Takeaway: Ionisation energy is a measure of how "greedy" an atom is for its electrons. The closer an electron is to a strong nucleus, the harder it is to remove.
4. Electronic Configuration
Electrons don't just swarm around the nucleus randomly. They live in specific quantum shells and orbitals.
Shells, Sub-shells, and Orbitals
Shells: Main energy levels (n = 1, 2, 3, 4).
Orbitals: A region of space where there is a high probability of finding an electron. Each orbital holds up to 2 electrons with opposite spins.
s-orbitals: Spherical shape.
p-orbitals: Dumbbell shape.
Memory Aid: "s" for Sphere, "p" for p-dumbbell!
Capacity of Shells
1st Shell: 2 electrons (one 1s orbital)
2nd Shell: 8 electrons (one 2s and three 2p orbitals)
3rd Shell: 18 electrons (one 3s, three 3p, and five 3d orbitals)
4th Shell: 32 electrons (one 4s, three 4p, five 4d, and seven 4f orbitals)
The Filling Rules
1. Aufbau Principle: Fill the lowest energy levels first.
2. Hund’s Rule: Electrons prefer to occupy orbitals alone before pairing up (like people on a bus!).
3. Pauli Exclusion Principle: Two electrons in the same orbital must have opposite spins (represented by up and down arrows).
Writing Configurations (Up to Z=36)
You write them like this: \( 1s^2 2s^2 2p^6... \)
Important Note: The 4s sub-shell is lower in energy than 3d, so it fills before 3d.
Example (Iron, Z=26): \( 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6 \)
Common Mistake: When atoms become ions, they lose electrons from the 4s sub-shell before the 3d sub-shell. Always remember: "First in, first out" for 4s!
Key Takeaway: Electronic configuration is the "address" of every electron in an atom. Use the 1s notation and remember the 4s/3d rule!
5. Periodicity
Periodicity is the repeating pattern of physical and chemical properties of elements as you move across the periods of the Periodic Table.
Trends in Period 2 and 3
Atomic Radii: Decreases across a period because the nuclear charge increases, pulling the shells closer.
Melting and Boiling Points: This depends on the structure and bonding:
1. Metals (Groups 1-3): Increase as the number of delocalised electrons increases (stronger metallic bonds).
2. Giant Covalent (Group 4): Very high (e.g., Diamond, Silicon) because many strong covalent bonds must be broken.
3. Simple Molecular (Groups 5-7): Low, as only weak London forces between molecules need to be overcome.
4. Noble Gases (Group 0): Very low, as they exist as individual atoms with very weak forces.
Evidence for Sub-shells: Small drops in IE between Group 2 and 3 (the electron is in a higher energy p-orbital) and Group 5 and 6 (electron-electron repulsion in a paired p-orbital) provide proof that sub-shells exist.
Key Takeaway: The Periodic Table is organized so that elements with similar electronic configurations (and therefore similar properties) fall into the same columns (Groups).
Final Quick Review Box
- Protons: Positive (+1), mass 1, in nucleus.
- Neutrons: Neutral (0), mass 1, in nucleus.
- Electrons: Negative (-1), mass ~0, in shells.
- Isotopes: Same protons, different neutrons.
- IE Factors: Nuclear charge, distance, shielding.
- Filling Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
- Group Jump: A huge jump in IE = a new shell is reached.