3.1.1 Atomic Structure (Physical Chemistry AS)
Hello future chemist! Atomic structure is the bedrock of chemistry. Everything we study—from bonding to reactions—depends on how atoms are built. Don't worry if this chapter seems dense; we're going to break it down into easy, bite-sized pieces!
In this section, you will learn about the tiny particles that make up atoms, how to count them, how chemists measure their mass with incredible accuracy, and the rules governing where electrons live.
3.1.1.1 Fundamental Particles: The Building Blocks
Every atom is made of three key subatomic particles. Understanding their properties is the essential starting point.
An atom consists of a central, dense nucleus (containing protons and neutrons) surrounded by orbiting electrons.
| Particle | Location | Relative Mass | Relative Charge |
|---|---|---|---|
| Proton | Nucleus | 1 | +1 |
| Neutron | Nucleus | 1 | 0 (Neutral) |
| Electron | Shells/Orbitals | \(\frac{1}{1836}\) (or negligible) | -1 |
Quick Key Facts:
- Atoms are electrically neutral because the number of protons (+1 charge) equals the number of electrons (-1 charge).
- Most of the mass of an atom is concentrated in the nucleus (protons and neutrons).
3.1.1.2 Mass Number and Isotopes
To identify an atom, we use two key numbers:
Atomic (Proton) Number (\(Z\))
The Atomic Number (\(Z\)) is the number of protons in the nucleus. This number defines the element. If the proton number changes, the element changes!
Memory Aid: Z is for P, as in Protons. P = Z.
Mass Number (\(A\))
The Mass Number (\(A\)) is the total number of protons and neutrons in the nucleus.
\[A = \text{Protons} + \text{Neutrons}\]
How to calculate the number of particles:
- Protons: Always equal to \(Z\).
- Electrons: In a neutral atom, electrons = protons. In an ion, you adjust for the charge:
- Positive ion (e.g., \(Na^{+}\)): Electrons = Protons - Charge.
- Negative ion (e.g., \(Cl^{-}\)): Electrons = Protons + Charge.
- Neutrons: Neutrons = \(A - Z\).
Example: A sodium ion (\(^{23}_{11}\text{Na}^{+}\))
\(Z=11\). Protons = 11.
\(A=23\). Neutrons = \(23 - 11 = 12\).
Charge is +1, meaning it lost one electron. Electrons = \(11 - 1 = 10\).
The Existence of Isotopes
Isotopes are atoms of the same element (meaning they have the same number of protons, \(Z\)) but different numbers of neutrons (meaning they have different mass numbers, \(A\)).
Example: Carbon-12 (\(^{12}\text{C}\)) has 6 neutrons, while Carbon-14 (\(^{14}\text{C}\)) has 8 neutrons. Both have 6 protons.
Because chemical reactions involve electrons, isotopes of an element have identical chemical properties (they react the same way), but slightly different physical properties (like density or mass).
Key Takeaway 1: Protons define the element (\(Z\)). Neutrons create isotopes. Electrons determine bonding and reactivity.
3.1.1.2 Mass Spectrometry: Weighing Atoms
How do we know atoms exist as isotopes? We use a powerful machine called a Time of Flight (TOF) Mass Spectrometer. This machine separates ions based on their mass-to-charge ratio (\(m/z\)).
The Principle of the TOF Mass Spectrometer
Imagine a race track for atomic particles. All racers must start with the same energy, but the lighter runners will finish faster.
- Ionisation: The sample is turned into positive ions. Electrons are knocked off the atoms/molecules (usually by electron impact or electrospray). This is essential because the particles must be charged so they can be accelerated by an electric field.
- Acceleration: The positive ions are accelerated by an electric field, giving all ions the same kinetic energy (KE).
- Ion Drift: The ions travel through a vacuum tube (the 'flight tube'). Since KE is constant, lighter ions travel faster than heavier ions. \[\text{Kinetic Energy} (KE) = \frac{1}{2}mv^2\] If KE is constant, a lower mass (\(m\)) must mean a higher velocity (\(v\)).
- Detection: When the ions hit the detector plate, they gain an electron, generating a current. The heavier ions take longer to reach the detector (longer time of flight).
- Data Analysis: The detector measures the time of flight for each ion and converts this into a mass-to-charge ratio (\(m/z\)) and registers the abundance (how many ions hit the detector).
Did you know? Modern mass spectrometers are so sensitive they can measure the mass of a large molecule (like a protein) with accuracy better than one part in a million!
Interpreting Simple Mass Spectra of Elements
A mass spectrum is a graph showing the relative abundance of ions plotted against their \(m/z\) ratio. For elements, the \(m/z\) ratio is usually equal to the relative isotopic mass.
- Each peak represents a specific isotope.
- The height of the peak shows the relative abundance (how common that isotope is).
Calculating Relative Atomic Mass (\(A_r\))
The Relative Atomic Mass (\(A_r\)) of an element is the weighted average mass of all its naturally occurring isotopes, relative to 1/12th the mass of a carbon-12 atom.
You must be able to calculate \(A_r\) from isotopic abundance data:
\[A_r = \frac{\sum (\text{Isotopic Mass} \times \text{Relative Abundance})}{\sum (\text{Relative Abundance})}\]
Example: If chlorine has 75% Cl-35 (mass 35.0) and 25% Cl-37 (mass 37.0).
\[A_r = \frac{(35.0 \times 75) + (37.0 \times 25)}{75 + 25} = \frac{2625 + 925}{100} = 35.5\]
We can also use mass spectrometry to determine the relative molecular mass (\(M_r\)) of a molecular compound, as the heaviest peak (the molecular ion peak) corresponds to the \(M_r\).
Key Takeaway 2: Mass spec identifies isotopes and their natural abundance, allowing us to calculate the weighted average relative atomic mass (\(A_r\)).
3.1.1.3 Electron Configuration
The arrangement of electrons is what gives an element its chemical personality. Electrons are arranged in shells, which are further divided into sub-shells (or orbitals).
Shells and Sub-shells
The principal quantum number, \(n\), defines the main energy shell (1, 2, 3...). Each shell contains different types of sub-shells:
- Shell 1 (\(n=1\)): Contains only the s sub-shell.
- Shell 2 (\(n=2\)): Contains s and p sub-shells.
- Shell 3 (\(n=3\)): Contains s, p, and d sub-shells.
- Shell 4 (\(n=4\)): Contains s, p, d, and f sub-shells.
We are required to know the configurations for atoms and ions up to \(Z=36\) (Krypton), focusing on the s, p, and d sub-shells.
Maximum Electrons per Sub-shell:
- s orbital: 2 electrons
- p orbitals (3 of them): 6 electrons
- d orbitals (5 of them): 10 electrons
The Filling Order (Aufbau Principle):
Electrons fill the lowest energy level first. The order is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ...
Note: The 4s sub-shell actually fills before the 3d sub-shell because it is slightly lower in energy.
Writing Electron Configurations (Example: Potassium, Z=19)
1. Count electrons (19).
2. Fill sub-shells sequentially:
\(1s^2 2s^2 2p^6 3s^2 3p^6 4s^1\)
Example: Iron ion (\(Fe^{3+}\), Z=26)
Neutral Fe: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\)
When forming positive ions, electrons are always removed from the highest principal quantum number (n) shell first. So, remove the 4s electrons before the 3d electrons.
\(Fe^{3+}\) (23 electrons): Remove 2 from 4s, and 1 from 3d.
\(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5\) (The 4s is now empty.)
Common Mistake Alert: Remember to always remove electrons from the 4s orbital *before* the 3d orbital when forming positive ions of transition metals!
Key Takeaway 3: Electron configuration follows the 1s, 2s, 2p, 3s, 3p, 4s, 3d order. When forming ions, 4s electrons are lost before 3d electrons.
3.1.1.3 Ionisation Energies
Ionisation energy provides the most direct evidence for the existence of electron shells and sub-shells.
Definition of First Ionisation Energy (\(IE_1\))
The First Ionisation Energy (\(IE_1\)) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
It is always an endothermic process (requires energy input).
Equation Writing: You must include state symbols \((g)\) and represent the removal of *one* electron.
\[\text{X}(g) \rightarrow \text{X}^{+}(g) + e^{-}\]
Successive Ionisation Energies
Successive Ionisation Energies (\(IE_2, IE_3\), etc.) measure the energy required to remove the second, third, and subsequent electrons.
Equation for the second IE:
\[\text{X}^{+}(g) \rightarrow \text{X}^{2+}(g) + e^{-}\]
Key Principle: Successive IEs always increase because you are removing an electron from an increasingly positive ion, meaning the remaining electrons are held tighter.
Evidence from Successive Ionisation Energies
When you plot successive IEs, you observe massive jumps in energy. These jumps occur when an electron is removed from a shell closer to the nucleus (or a full, stable sub-shell).
- Electrons removed from a shell closer to the nucleus feel a much stronger electrostatic attraction.
- The number of electrons removed before the first big jump tells you how many electrons are in the outermost shell (valence electrons).
Analogy: Removing electrons is like peeling an onion. The first layers come off easily, but once you hit the next layer, it takes significantly more effort (energy).
Trends in First Ionisation Energy (IE)
The size of the first IE depends on three factors (all related to the attraction between the nucleus and the outer electron):
- Nuclear Charge (\(Z\)): More protons = stronger attraction = higher IE.
- Atomic Radius: Larger atom = valence electron is further away = lower IE.
- Shielding: More inner electrons blocking the nuclear charge = less attraction for outer electron = lower IE.
Using IE to Provide Evidence for Sub-shells (Period 3 and Group 2)
The general trend across a period (like Period 3: Na to Ar) is that IE increases, reflecting the increasing nuclear charge while shielding remains relatively constant (same main shell). However, there are two important dips:
1. Dip between Group 2 and Group 13 (e.g., Mg to Al in Period 3):
- The electron being removed from Al (\(3s^2 3p^1\)) is from the 3p sub-shell, while for Mg (\(3s^2\)), it is from the 3s sub-shell.
- The 3p electron is at a slightly higher energy level and is shielded by the full 3s sub-shell.
- Result: Al's first IE is slightly lower than Mg's. (Evidence for the start of the p sub-shell).
2. Dip between Group 15 and Group 16 (e.g., P to S in Period 3):
- The electron being removed from P (\(3p^3\)) is from an orbital that is half-filled (stable).
- The electron being removed from S (\(3p^4\)) is from an orbital containing a pair of electrons.
- The mutual repulsion between the paired electrons in the S orbital makes it easier to remove one of them (less energy required).
- Result: S's first IE is slightly lower than P's. (Evidence for electron pairing/sub-orbital stability).
Group 2 Trend (Be to Ba):
As you go down Group 2, the first IE decreases significantly.
- Reason: Although the nuclear charge increases, the outer electron is in a new shell further from the nucleus (larger atomic radius).
- The increasing number of inner shells provides much greater shielding.
- The effect of increased shielding and distance outweighs the increased nuclear charge, making the outer electron easier to remove. (Evidence for the addition of new, larger shells).
Key Takeaway 4: Ionisation energies confirm the shell structure (large jumps) and the sub-shell structure (small dips caused by shielding and electron repulsion).