Atomic Structure and Electron Configuration Notes

Fundamental Particles

Plum Pudding Model:
- Early atomic model. Atoms were thought to consist of a sphere of positive charge, with small negative charges (electrons) distributed evenly within it.
Electron Shell Model:
- Current model.
- The atom consists of a small, dense, positively charged central nucleus surrounded by orbiting electrons in electron shells.
- Discovered in the Rutherford scattering experiment in 1911.
Nucleus:
- Composed of protons and neutrons.
- Carries an overall positive charge.
- Contains almost the entire mass of the atom.
Neutral Atom:
- The number of electrons is equal to the number of protons, due to charge balance.
Fundamental Particles:
- Proton:
  - Relative Charge: $+1$
  - Relative Mass: $1$
- Neutron:
  - Relative Charge: $0$
  - Relative Mass: $1$
- Electron:
  - Relative Charge: $-1$
  - Relative Mass: $1/1840$
Electron Shell Capacity:
- The maximum number of electrons that can be held by a single shell depends on the shell number ( $n$ ).
- Calculated using the formula: $2n^2$
- Example: Electrons in shell 2 = $2(2^2) = 8$ electrons
Filling Order:
- Each electron shell must fill before the next one can hold any electrons.

Mass Number and Isotopes

Mass Number (A):
- Represented by $A$ .
- Calculated as the sum of protons and neutrons in an atom.
Atomic Number (Z):
- Represented by $Z$ .
- Equal to the number of protons in an atom.
Calculating Particle Quantity:
- Using $A$ and $Z$ , the number of protons, neutrons, and electrons can be determined.
- Example:
  - Atomic number = 7
  - Mass number = 14
  - Proton number = 7
  - Neutron number = $14 - 7 = 7$
Relative Atomic Mass (Ar):
- Defined as the mean mass of an atom of an element, divided by one-twelfth of the mean mass of an atom of the carbon-12 isotope.
- Takes into account the relative abundances of different isotopes of an element.
Isotopes:
- Atoms of the same element with the same atomic number ( $Z$ ), but a different number of neutrons, resulting in a different mass number ( $A$ ).
Chemical Properties of Isotopes:
- Neutral atoms of isotopes react chemically in the same way because they have the same proton number and electron configuration.
- The sharing and transfer of electrons is unaffected.
Physical Properties of Isotopes:
- Different mass numbers mean they have different physical properties.
Isotope Examples (Hydrogen):
- Hydrogen: 1 proton, 0 neutrons
- Deuterium: 1 proton, 1 neutron
- Tritium: 1 proton, 2 neutrons
Ions:
- Formed when an atom loses or gains electrons.
- Results in an overall charge (no longer neutral).

Mass Spectrometry

Purpose:
- An analytical technique used to identify different isotopes and find the overall relative atomic mass of an element.
Time of Flight (TOF) Mass Spectrometry:
- Records the time it takes for ions of each isotope to reach a detector.
- Produces spectra showing each isotope present.
Steps in TOF Mass Spectrometry:
1. Ionization:
  - A sample of an element is vaporized and injected into the mass spectrometer.
  - A high voltage is passed over the chamber.
  - This causes electrons to be removed from the atoms, leaving +1 charged ions in the chamber.
2. Acceleration:
  - Positively charged ions are accelerated towards a negatively charged detection plate.
3. Ion Drift:
  - Ions are deflected by a magnetic field into a curved path.
  - The radius of their path depends on the charge and mass of the ion.
4. Detection:
  - When positive ions hit the negatively charged detection plate, they gain an electron, producing a flow of charge.
  - The greater the abundance, the greater the current produced.
5. Analysis:
  - Current values are used in combination with the flight times to produce a spectra print-out.
  - The printout displays the relative abundance of each isotope.
2+ Charged Ions:
- During ionization, a 2+ charged ion may be produced.
- This means it will be affected more by the magnetic field, producing a curved path of smaller radius.
- As a result, its mass to charge ratio ( $m/z$ ) is halved and this can be seen on spectra as a trace at half the expected $m/z$ value.
Calculating Ar from Spectra:
- Example: If a spectra shows two isotopes with mass 10 and 12, with abundances of 75% and 25% respectively:
- $A_r = \frac{(10 \times 75) + (12 \times 25)}{(75 + 25)} = 10.5$
Chlorine Spectra:
- Spectra produced by the mass spectrometry of chlorine display a characteristic pattern.
- A 3:1 ratio for $Cl^+$ ions and a 3:6:9 ratio for $Cl_2^+$ ions.
- This is because one isotope is more common than the other and the chlorine molecule can form in different combinations.
- Examples of $Cl_2^+$ combinations:
  - $^{70}Cl_2^+ = ^{35}Cl + ^{35}Cl$
  - $^{72}Cl_2^+ = ^{35}Cl + ^{37}Cl \text{ OR } = ^{37}Cl + ^{35}Cl$
  - $^{74}Cl_2^+ = ^{37}Cl + ^{37}Cl$

Electron Configuration

Electron Orbitals:
- Electrons are held in clouds of negative charge called orbitals.
- Different types of orbitals: s, p, d, and f.
- Each has a different shape.
Orbital Correspondence to Periodic Table Blocks:
- Each element in a block has outer electrons in that orbital.
Electron Capacity of Orbitals:
- s-orbital = 2 electrons
- p-orbital = 6 electrons
- d-orbital = 10 electrons
Orbital Filling Order:
- The energy of the orbitals increases from s to d meaning the orbitals are filled in this order.
- Each orbital is filled before the next one is used to hold electrons.
- Example: Sodium (Na) has 11 electrons:
  - $Na = 1s^22s^22p^63s^1$
  - It has 3 energy levels and 4 orbitals holding the 11 electrons.
Electron Spin:
- Within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible.
- Electrons in the same orbital must have opposite spin.
- Spin is represented by arrows.
Rules for Writing Electron Configurations:
1. The lowest energy orbital is filled first.
2. Electrons with the same spin fill up an orbital first before pairing begins.
3. No single orbital holds more than 2 electrons.

Exceptions to the Rules

Instability due to Unpaired Electron Spins:
- If electron spins are unpaired and therefore unbalanced, it produces a natural repulsion between the electrons making the atom very unstable.
- If this is the case, the electrons may take on a different arrangement to improve stability.
Example of Electron Configuration Adjustment:
- The $3p^4$ orbital contains a single pair of electron with opposite spin making it unstable.
- Therefore the electron configuration changes to become $3p^34s^1$ which is a much more stable arrangement.

Ionisation Energy

Definition:
- The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state.
- Measured in $kJmol^{-1}$ .
- Na(g) ---------> Na^+(g) + e^-
Successive Ionisation Energies:
- Occur when further electrons are removed.
- Usually requires more energy because as electrons are removed the electrostatic force of attraction between the positive nucleus and the negative outer electron increases.
- More energy is therefore needed to overcome this attraction so ionisation energy increases.
Trends in First Ionisation Energy:
- Along a Period: First ionisation energy increases due to a decreasing atomic radius and greater electrostatic forces of attraction.
- Down a Group: First ionisation energy decreases due to an increasing atomic radius and shielding which reduces the effect of the electrostatic forces of attraction.
Evidence for Atomic Orbital Theory:
- When successive ionisation energies are plotted on a graph, a sudden large increase indicates a change in energy level.
- This is because the electron is being removed from an orbital closer to the nucleus so more energy is required to do so.
Anomalies in Ionisation Energy:
- The first ionisation energy of Aluminium is lower than expected due to a single pair of electrons with opposite spin.
- As a result there is a natural repulsion which reduces the amount of energy needed to be put in to remove the outer electron.