Atomic Structure

Question 1

Relative charges and relative masses.

Answer 1

P: +1, 1
N: 0, 1
E: 1/1840

Question 2

Behaviour in electric field

Answer 2

• Protons: Deflected towards negative potential
  
  • Electrons: Deflected towards positive potential
  • Neutrons: Not deflected at all
  • Size of deflection: Depends on the relative masses of the particles –> The lighter the particles the more the deflection –> e.g. Electrons are lighter than protons and so are deflected more than protons
  • Magnitude of deflection is proportional to charge/mass ratio (numerically, charge=proton number; mass=mass number)

Question 3

Isotopes

Answer 3

• Atoms of the same element that have the same number of protons but different number of neutrons
  
  • Have the same chemical properties but different physical properties

Question 4

Atoms or ions are called…

Answer 4

• Isoelectronic: contain same number of electrons
  
  • Isotonic: contain same number of neutrons
  • Isotopic: contain same number of protons

Question 5

First I.E.

Answer 5

• The energy required to remove one mole of the outermost electrons from one mole of gaseous M atoms to form one mole of gaseous M+ cations
  
  • M(g) –> M+(g) + e-

Question 6

Second I.E.

Answer 6

• The energy required to remove one mole of the outermost electrons from one mole of gaseous M+ cations to form one mole of gaseous M2+ cations
  
  • M+(g) –> M2+(g) + e-

Question 7

Principal quantum shells

Answer 7

• The closer the quantum shell is to the nucleus, the greater the attraction between the electrons occupying it and the nucleus, the lower the energy of the electrons

Question 8

Atomic orbitals

Answer 8

• An orbital is defined as the region of space within which there is a 95% probability of locating a particular electron in an atom
  
  • Each orbital has a characteristic energy level and shape
  • Orbitals in the same sub-shell are equivalent in energy but they differ in their orientation in space

Question 9

3 rules

Answer 9

• Aufbau Principle: Electrons are placed in the orbital of the lowest energy, then the orbital of the next lowest energy and so on (fill in 4s before 3d)
  
  • Pauli Exclusion Principle: An orbital can only accommodate a maximum of 2 electrons, which must have opposite spins
  • Hund’s Rule: When electrons are added to the same subshell, electrons occupy degenerate orbitals singly and with the same spin before pairing occur. This minimises inter-electronic repulsion.

Question 10

Periodic relationships

Answer 10

• Period number: Number of principal quantum shell of the element
  
  • Group number: Number of valence electrons

Question 11

Factors influencing magnitude of ionisation energy

Answer 11

1. Size of positive nuclear charge (attractive forces) and shielding effect of inner electrons (repulsive forces): Effective Nuclear Charge
   (i) Nuclear Charge (attractive force)
   
   • The attractive forces exerted by the nucleus on the valence electrons in an atom are dependent on the number of protons in the nucleus
   • As the number of protons increases, nuclear charge of the particle gets more and more positive and hence the attraction of the nucleus for the electron increases and consequently, the ionisation energy increases.
     (ii) Shielding Effect (repulsive force)
   • The presence of electrons in the inner shells between the outermost electron and the nucleus shields the outermost electron from the attraction from the nucleus and consequently reduces the attraction between them
   • The greater the number of inner shell electrons, the greater the shielding effect and consequently, the ionisation energy decreases
   • Electrons within the same shell exert a negligible shielding effect on each other
     (iii) Effective Nuclear Charge (ENC = NC - SE)
   • Higher ENC –> stronger attraction of the valence electrons to the nucleus –> higher amount of energy to remove it
   2. The orbital where the electron comes from
      - The further the average distance from the nucleus of the orbital is, the lower the amount of energy is required to remove the electron in it
      - This factor can only be used if the electrons removed come from the same principal quantum shell but different sub-shells
   3. Inter-electronic repulsion between electrons in the same orbital
      - Ionisation energy can be lowered if the electron is being removed from an orbital that is paired as compared to the case where the orbital contains a singly unpaired electron.
      - In the case of unpaired electron, mutual repulsion between the electrons is minimised while in the case of paired electron, the two electrons experience inter-electronic repulsion and hence more easily removed
   4. The charge of cation
      - As the charge of the cation increases –> a greater amount of energy is required to remove the subsequent electron due to an increase in attraction by the same number of protons on a fewer number of electrons –> results in increase in the ENC

Question 12

Trends in I.E ACROSS a period

Answer 12

• Observation 1: First IE of the elements generally increases across a period
  –> cuz, across a period –> nuclear charge of the elements increases as the number of protons increases
  –> shielding effect remains approximately constant as the elements have the same number of inner electronic shells of electrons
  –> effective nuclear charge increases
  –> the amount of energy required to remove the valence electron increases
  
  • Observation 2: First IE of a Group 13 is lower than expected as compared to that of a Group 2 element
    –> cuz, the average distance from the nucleus of a np orbital is slightly larger than that of a ns orbital, so less energy required to remove the np electron
  • Observation 3: First IE of a Group 16 element is lower than expected as compared to that of a Group 15 element
    –> the electron to be removed is one of three 2p electrons and it is unpaired and the mutual repulsion between these electrons is minimised
    –> however in the other electron to be removed in the p-orbital is paired and since these to electrons experience inter-electronic repulsion, it is easier to remove one of them than p electron
    –> hence less energy required to remove it

Question 13

Trends in I.E DOWN a group

Answer 13

• Observation 1: First IE of the elements generally decreases down a group
  –> nuclear charge of the elements increases due to an increase in the number of protons the elements have
  –> shielding effect also increases cuz of the increase in the number of electronic shells
  –> increase in shielding effect outweighs the increase in the nuclear charge resulting in a decrease in ENC
  –> amount of energy required to remove the valence electron decreases

Question 14

Summary

Answer 14

group 1 vs 2: nuclear charge
group 2 vs 13: ns vs np
group 15 vs 16: interelectronic repulsion
group 16: down group, increase in screening effect