Welcome to Atomic Structure and the Periodic Table!
Welcome to the first step of your AS Level Chemistry journey! This chapter is the foundation of everything else you’ll learn. We are going to explore what atoms are made of, how we can "weigh" them, and why the Periodic Table is organized the way it is. Don't worry if some of the math or the patterns seem a bit strange at first—once you see the logic behind them, it all starts to click like a puzzle!
1. The Heart of the Atom: Sub-atomic Particles
Atoms are the tiny building blocks of everything around you. Even though they are incredibly small, they are made of three even smaller pieces: protons, neutrons, and electrons.
Mass and Charge
Think of the nucleus (the center) as a heavy suitcase and the electrons as tiny flies buzzing around it.
- Protons: Found in the nucleus. Relative mass = 1. Relative charge = +1.
- Neutrons: Found in the nucleus. Relative mass = 1. Relative charge = 0 (they are neutral!).
- Electrons: Found in shells orbiting the nucleus. Relative mass = 1/1840 (so small we often say "negligible"). Relative charge = -1.
Atomic Number and Mass Number
To identify an atom, we look at its numbers on the Periodic Table:
1. Atomic (proton) number (Z): This is the number of protons in the nucleus. It defines the element! If you change the proton number, you change the element.
2. Mass number (A): This is the total number of protons + neutrons.
Quick Calculation Trick:
To find the number of neutrons, just subtract the bottom number from the top number: \( \text{Neutrons} = \text{Mass Number} - \text{Atomic Number} \).
In a neutral atom, the number of electrons is always equal to the number of protons.
Key Takeaway: Protons and neutrons live in the center (nucleus) and provide the mass; electrons orbit on the outside and provide the charge.
2. Isotopes and Atomic Mass
Nature isn't always identical. Sometimes, atoms of the same element have different numbers of neutrons. We call these isotopes.
Isotopes are atoms of the same element with the same number of protons but a different number of neutrons. This means they have the same atomic number but different mass numbers.
The Carbon-12 Scale
Because atoms are too small to weigh on a normal scale, we compare them to a standard: the Carbon-12 isotope.
- 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 (RAM or \(A_r\)): The weighted mean mass of an atom of an element compared with 1/12th of the mass of an atom of carbon-12.
Analogy: If you have a bag of marbles where 75% weigh 10g and 25% weigh 12g, the "weighted average" is the RAM.
Key Takeaway: Isotopes react chemically the same way because they have the same number of electrons, but they have different physical masses.
3. Mass Spectrometry: Weighing Atoms
How do we actually know how many isotopes an element has? We use a machine called a Mass Spectrometer.
Calculating Relative Atomic Mass
You might be asked to calculate the RAM from a graph (a mass spectrum). The x-axis is \(m/z\) (mass/charge ratio) and the y-axis is abundance (percentage).
\( RAM = \frac{\sum (\text{isotopic mass} \times \text{relative abundance})}{\text{total abundance}} \)
Diatomic Molecules (like Chlorine, \(Cl_2\))
Chlorine has two main isotopes: \(^{35}Cl\) and \(^{37}Cl\). When they form a molecule (\(Cl_2\)), you can get three different combinations: 35+35 (mass 70), 35+37 (mass 72), and 37+37 (mass 74). Chemistry exams love asking you to predict the height of these peaks based on probability!
The Molecular Ion Peak
For a whole molecule, the mass spectrometer gives a peak called the molecular ion peak (\(M^+\)). This peak tells you the relative molecular mass of the substance.
Key Takeaway: The mass spectrometer is like a very sensitive set of scales that separates particles by their mass.
4. Ionisation Energy: The Cost of an Electron
If you want to pull an electron away from an atom, you have to "pay" with energy. This is Ionisation Energy (IE).
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 IE:
1. Nuclear Charge: More protons = more "pull" on electrons (IE increases).
2. Shielding: More inner shells of electrons "block" the pull of the nucleus (IE decreases).
3. Distance (Atomic Radius): Further away the electron is, the easier it is to remove (IE decreases).
Trends to Remember:
- Down a Group: IE decreases. Even though there are more protons, the extra shells increase shielding and distance, making the outer electron easier to steal.
- Across a Period: IE generally increases. The nuclear charge increases (more protons) but the shielding stays roughly the same.
Did you know? There are small "dips" in the trend across a period. These happen because of sub-shells (like moving from an \(s\) sub-shell to a \(p\) sub-shell), which provides evidence that electrons aren't just in big shells, but in specific sub-layers!
Key Takeaway: Successive ionisation energies (removing 1st, 2nd, 3rd electron) always increase, and big "jumps" in energy tell us when we've moved to a new, inner shell closer to the nucleus.
5. Electronic Configuration: Where are the electrons?
Electrons don't just fly around randomly. They live in quantum shells, sub-shells, and orbitals.
Shells and Sub-shells
- Quantum Shells: Numbered 1, 2, 3, 4. The max electrons they can hold are calculated by \( 2n^2 \). (Shell 1 = 2, Shell 2 = 8, Shell 3 = 18, Shell 4 = 32).
- Sub-shells: Shells are split into \(s, p,\) and \(d\) sub-shells.
- \(s\)-subshell: 1 orbital (max 2 electrons)
- \(p\)-subshell: 3 orbitals (max 6 electrons)
- \(d\)-subshell: 5 orbitals (max 10 electrons)
Orbitals
An orbital is a region of space where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons, and they must have opposite spins.
- s-orbital: Spherical shape.
- p-orbital: Dumbbell shape (there are three of these: \(p_x, p_y, p_z\)).
Filling Rules
1. Hund's Rule: Electrons fill orbitals singly before pairing up (like people getting on a bus—everyone wants their own seat first!).
2. Pauli Exclusion Principle: Two electrons in the same orbital must have opposite spins (represented by up and down arrows).
Writing Configurations
You need to know how to write these for atoms up to \(Z=36\) (Krypton).
Example for Magnesium (12 electrons): \( 1s^2 2s^2 2p^6 3s^2 \)
Key Takeaway: Electronic configuration determines how an atom reacts. Elements in the same group have the same outer shell configuration!
6. Periodicity: Patterns in the Table
Periodicity is the repeating pattern of physical and chemical properties as you go across different periods.
Blocks
The table is divided into blocks based on which sub-shell the highest-energy electron is in:
- s-block: Groups 1 and 2.
- p-block: Groups 3 to 0.
- d-block: The transition metals in the middle.
Melting and Boiling Points
The trends in melting points across Periods 2 and 3 depend on bonding and structure:
- Metals (Li, Be / Na, Mg, Al): Boiling points increase as the "sea of delocalised electrons" gets stronger.
- Giant Covalent (C / Si): Very high boiling points because you have to break many strong covalent bonds.
- Simple Molecular (N, O, F, Ne / P, S, Cl, Ar): Low boiling points because you only have to break weak intermolecular forces.
Quick Review Box:
- Atomic Radii: Get smaller across a period (stronger pull from nucleus).
- Ionisation Energy: Generally increases across a period.
- Chemical Properties: Determined by the outer shell electrons.
Key Takeaway: The Periodic Table isn't just a list; it's a map. If you know where an element is, you can predict how it will behave!
Don't worry if this seems like a lot to memorize! Practice writing out electronic configurations and drawing the IE graphs, and you'll find the patterns start to repeat themselves everywhere in Chemistry.