Exam Study Notes

Fundamental Particles

• Atomic structure model has evolved.
• Current model: small, dense, central nucleus surrounded by orbiting electrons in electron shells.
• Rutherford scattering experiment (1911) led to this discovery.
• Nucleus: protons and neutrons (positive charge).
• Almost entire mass of the atom is in the nucleus.
• Neutral atom: number of electrons = number of protons.

Particle Properties

• Proton:

  • Relative Charge: +1
  • Relative Mass: 1

• Neutron:

  • Relative Charge: 0
  • Relative Mass: 1

• Electron:

  • Relative Charge: -1
  • Relative Mass: 1/1840

• Maximum number of electrons in a shell depends on shell number:

  • Formula: $2n^2$ (where n is the shell number)
  • Example: Electrons in shell 2 = $2(2^2) = 8$

• Each electron shell must fill before the next one.

Atomic Number and Mass Number

• Mass number:
  • Represented by A
  • Calculation: sum of protons and neutrons in an atom.
• Atomic number:
  • Represented by Z
  • Equal to the number of protons.
  • Also referred to as proton number.

Calculations

• Using these numbers, the quantity of each fundamental particle in an atom can be calculated.

Example

• Atomic number = 7
• Mass number = 14
• Proton number = 7
• Neutron number = 14 - 7 = 7

Isotopes

• Atoms of the same element.
• Same atomic number.
• Different number of neutrons, resulting in a different mass number.
• Neutral atoms of isotopes react chemically the same way (same proton number and electron configuration).
• Sharing and transfer of electrons is unaffected.
• Different mass numbers mean different physical properties.

Examples

• Hydrogen:
  • 1 proton, 0 neutrons
• Deuterium:
  • 1 proton, 1 neutron
• Tritium:
  • 1 proton, 2 neutrons

Relative Masses

Relative Atomic Mass (Ar)

• Definition: The mean mass of an atom of an element, relative to 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.

Relative Isotopic Mass

• Definition: The isotopic mass of an isotope relative to one-twelfth of the mean mass of an atom of the carbon-12 isotope.

Relative Molecular Mass (Mr)

• Definition: The mean mass of a molecule of a compound, relative to one-twelfth of the mean mass of an atom of the carbon-12 isotope.
• Calculation: Add together the separate Ar values of the component elements.

Relative Formula Mass

• Similar to Mr but is used for compounds with giant structures.

Ions and Mass Spectrometry

• Ions: formed when an atom loses or gains electrons.
• No longer neutral; have an overall charge.
• Useful in mass spectrometry.
• Identifies different isotopes and finds overall relative atomic mass of an element.

Time of Flight (TOF) Mass Spectrometry

• Records time it takes for ions of each isotope to reach a detector.
• Produces spectra showing each isotope present.

Steps:

1. Ionisation:
   • Sample is vaporized and injected into the mass spectrometer.
   • High voltage passed over the chamber.
   • Causes electrons to be removed from atoms (ionized), leaving +1 charged ions.
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.
   • Radius depends on charge and mass of the ion.
4. Detection:
   • Positive ions hit negatively charged detection plate.
   • Gain an electron, producing a flow of charge.
   • Greater abundance = greater current produced.
5. Analysis:
   • Current values used with flight times.
   • Produces spectra print-out with relative abundance of each isotope displayed.

Additional Notes

• During ionization, a 2+ charged ion may be produced.
• Affected more by the magnetic field, producing a curved path of smaller radius.
• Mass to charge ratio (m/z) is halved.
• Can be seen on spectra at half the expected m/z value.

Calculation of Ar

• Multiply each m/z value by its abundance.
• Add each of these together.
• Divide by the total abundance of all species present.
• $Ar = \frac{\sum(m/z \times abundance)}{\sum abundance}$

Example

• From a spectra:
  $Ar = \frac{(10 \times 75) + (12 \times 25)}{(75 + 25)} = 10.5$
• Using calculated Ar value, the element can be identified by referring to the Periodic Table.
• Tallest peak on a mass spectrum corresponds to the relative molecular mass of the molecule.
• This peak is known as the molecular ion peak and is formed from the $M^+$ species.

Chlorine Spectra

• Spectra produced by mass spectrometry of chlorine display a characteristic pattern.
• 3:1 ratio for $Cl^+$ ions.
• 3:6:9 ratio for $Cl_2^+$ ions.
• One isotope is more common than the other.
• Chlorine molecule can form in different combinations.

Example:

• $^{70}Cl_2^+ = ^{35}Cl + ^{35}Cl$
• $^{72}Cl_2^+ = ^{35}Cl + ^{37}Cl OR ^{37}Cl + ^{35}Cl$
• $^{74}Cl_2^+ = ^{37}Cl + ^{37}Cl$

Ionisation Energy

• Definition: The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state.
• Measured in $kJmol^{-1}$.
• $Na(g) \longrightarrow Na^+(g) + e^-$
• Successive ionisation energies occur when further electrons are removed.
• 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 ionization energy increases.
• First ionisation energy follows trends within the Periodic Table.
• Influenced by proton-electron forces of attraction and electron shielding.

Trends:

• Along a Period:
  • First ionisation energy increases.
  • Due to decreasing atomic radius and greater electrostatic forces of attraction.
• Down a Group:
  • First ionisation energy decreases.
  • Due to increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction.
• 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.
• This large energy increase provides supporting evidence for the atomic orbital theory.
• 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.

Electron Configurations

• Scientific ideas on electronic configurations have developed over time.
• The current accepted model is based on the following evidence:

Evidence:

1. Emission spectra provide evidence for the existence of quantum shells.
2. Successive ionisation energies provide evidence for quantum shells within atoms and suggest the group to which the element belongs.
3. The first ionisation energy of successive elements provides evidence for electron subshells.

Electron Orbitals

• Electrons are held in clouds of negative charge called orbitals.
• Different types: s, p, d, and f.
• Each can hold up to two electrons with opposite spins.
• Each has a different shape.
• These orbitals correspond with blocks on the Periodic Table.
• Each element in the block has outer electrons in that orbital.
• Each subshell has a different number of orbitals and therefore can hold a different number of electrons before the next one is filled:

Subshell Capacities:

• s-subshell = 2 electrons
• p-subshell = 6 electrons
• d-subshell = 10 electrons
• 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

• Has 11 electrons.
• Configuration: $Na = 1s^22s^22p^63s^1$
• Has 3 energy levels and 4 orbitals holding the 11 electrons.

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 spins.
• Spin is represented by opposite 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

• 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

• The $3p^4$ orbital contains a single pair of electrons with opposite spins, making it unstable:
• Therefore, the electron configuration changes to become $3p^34s^1$, which is a much more stable arrangement.

Periodicity

• Periodicity refers to the study of patterns of physical, atomic, and chemical properties within the Periodic Table that repeat regularly.
• The Periodic Table arranges the known elements according to their proton number.
• All the elements along a period have the same number of electron shells.
• All the elements down a group have the same number of outer electrons; this number is indicated by the group number.

Element Blocks

• Elements are classified into blocks within the Periodic Table that show electron configuration:
  • s-block = groups 1 and 2
  • p-block = groups 3 to 0
  • d-block = transition metals
  • f-block = radioactive elements
• These different electron configurations are often linked to other trends within the Periodic Table.
• Periodicity is the study of these trends.

Atomic Radius

• Along a period:

  • Atomic radius decreases.
  • Due to an increased nuclear charge for the same number of electron shells.
  • The outer electrons are pulled in closer to the nucleus as the increased charge produces a greater attraction.
  • As a result, the atomic radius for that element is reduced.

• Down a group:

  • Atomic radius increases.
  • With each increment down a group, an electron shell is added each time.
  • This increases the distance between the outer electrons and the nucleus, reducing the power of attraction.
  • More shells also increase electron shielding, where the inner shells create a ‘barrier’ that blocks the attractive forces.
  • Therefore, the nuclear attraction is reduced further, and atomic radius increases.

Ionisation Energy

• Along a period:

  • Ionisation energy increases.
  • The decreasing atomic radius and increasing nuclear charge means that the outer electrons are held more strongly, and therefore more energy is required to remove the outer electron and ionise the atom.

• Down a group:

  • Ionisation energy decreases.
  • The nuclear attraction between the nucleus and outer electrons reduces, and increasing amounts of shielding means less energy is required to remove the outer electron.

Physical Properties of Period 2

Melting Points

• The melting points of the period two elements peak towards the middle of the period due to the different bond strength and structures:
• Lithium and Beryllium:
  • Have metallic bonding.
  • Their melting points increase due to greater positive charged ions (Li = +1, Be = +2).
  • This also means more electrons are released as free electrons in the Beryllium lattice, so the attractive electrostatic forces are greater than for Lithium.
• Boron and Carbon:
  • From giant covalent lattices with very strong covalent bonds in up to three dimensions.
  • These covalent bonds require a lot of energy to break, giving them very high melting points.
• Nitrogen, Oxygen, Fluorine and Neon:
  • Are all small, simple covalent molecules held with weak van der Waals forces.
  • These intermolecular forces don’t require much energy to overcome, so these molecules have relatively low, similar melting points.

Ionisation Energies

• First ionisation energies follow a general increasing trend along period 2.
• This is due to the decreasing atomic radius and increasing nuclear charge, so outer electrons are held more strongly.
• Boron and Oxygen are exceptions to this trend due to the quantum behaviour of the electrons.
• The electron configurations of these elements contain unpaired electrons that require less energy to remove, resulting in a lower first ionisation energy.

Physical Properties of Period 3

Melting Points

• The melting points of the period three elements is linked to the bond strength and structure:

• Sodium, Magnesium and Aluminium:

• Are all metals with metallic bonding.

• Their melting points increase due to greater positive charged ions (Na = +1, Mg = +2, Al = +3).

• This also means more electrons are released as free electrons, so the attractive electrostatic forces increase from Na to Al.

• Silicon:

• Is macromolecular, meaning it has a very strong covalent structure.

• These covalent bonds require a lot of energy to break, giving it a very high melting point.

• Phosphorus, Sulphur and Chlorine:

• Are all simple covalent molecules held with weak van der Waals forces.

• These intermolecular forces don’t require much energy to overcome, so these molecules have relatively low, similar melting points.

• Argon:

• Is a noble gas that exists as individual atoms with a full outer shell of electrons.

• This makes the atom very stable and the van der Waals forces between them very weak.

• As a result, the melting point of Argon is very low and it exists as a gas at room temperature.

Ionisation Energies

• First ionisation energies follow a general increasing trend along period 3.
• This is due to the decreasing atomic radius and increasing nuclear charge, so outer electrons are held more strongly.
• Aluminium and Sulfur are exceptions to this trend due to the quantum behaviour of the electrons.
• The electron configurations of these elements contain unpaired electrons that require less energy to remove, resulting in a lower first ionisation energy.