Lecture 1: Atomic Structure, Masses, and Ionization Energy

Basic Structure of Atoms

  • Atoms consist of three sub-atomic particles:

    • Protons: have a relative mass of 1 and a relative charge of +1.

    • Neutrons: have a relative mass of 1 and a relative charge of 0.

    • Electrons: have a relative mass of 0 and a relative charge of -1.

Isotopes

  • Definition: Atoms of the same element with a different number of neutrons.

  • Example: Isotopes of Carbon

    • 12C^{12}C: 6 protons, 6 neutrons

    • 13C^{13}C: 6 protons, 7 neutrons

    • 14C^{14}C: 6 protons, 8 neutrons

Ions

  • Definition: Atoms that have gained or lost electrons (different number of protons/electrons).

  • Examples:

    • Bromine (Br):

    • Neutral Br: 35 protons, 45 neutrons, 35 electrons.

    • Anion Br-: 36 electrons.

    • Magnesium (Mg):

    • Neutral Mg: 12 protons, 12 neutrons, 12 electrons.

    • Cation Mg2+: 10 electrons.

    • Copper (Cu):

    • Neutral Cu: 29 protons, 36 neutrons, 29 electrons.

    • Cation Cu1+: 28 electrons,

    • Cation Cu2+: 27 electrons.

Relative Masses

  • Relative Atomic Mass (Ar): average mass of an atom compared to 1/12 of the mass of a 12C^{12}C atom.

  • Relative Isotopic Mass: mass of an atom of an isotope compared to 1/12 the mass of a 12C^{12}C atom.

  • Relative Molecular Mass (Mr): average mass of a molecule compared to 1/12 of the mass of a 12C^{12}C atom.

Calculating Relative Atomic Mass, Ar

  • Mass spectrometer used for finding masses of different isotopes.

  • Example calculation:

    1. Relative abundance of Carbon isotopes:

    • 98.93×12=1187.1698.93 \times 12 = 1187.16

    • 1.06×13=13.781.06 \times 13 = 13.78

    • 0.01×14=0.140.01 \times 14 = 0.14

    1. Totalling the values:

    • 1187.16+13.78+0.14=1201.081187.16 + 13.78 + 0.14 = 1201.08

    1. Dividing by 100:

    • 1201.08100=12.0108\frac{1201.08}{100} = 12.0108

Relative Molecular Mass, Mr

  • Calculation consists of summing the atomic masses of all elements in the molecule:

    • Water (H2O):

    • O = 16.0, H = 1.0

    • Mr=16.0+1.0+1.0=18.0g/molMr = 16.0 + 1.0 + 1.0 = 18.0 g/mol

    • Carbon Dioxide (CO2):

    • C = 12.0, O = 16.0

    • Mr=12.0+16.0+16.0=44g/molMr = 12.0 + 16.0 + 16.0 = 44 g/mol

Electronic Structure

  • Electrons have fixed energies and move around the nucleus in shells (energy levels). Each shell can hold a different number of electrons:

    • 1st Shell (n=1): 2 electrons

    • 2nd Shell (n=2): 8 electrons

    • 3rd Shell (n=3): 18 electrons

    • 4th Shell (n=4): 32 electrons

Ionisation Energies

  • Definition: Energy required to remove electrons from atoms, forming ions.

  • First Ionisation Energy (IE1): Energy to remove one electron:

    • X(g)→X1+(g)+e−X(g) \rightarrow X^{1+}(g) + e^-

  • Second Ionisation Energy (IE2): Energy to remove a second electron:

    • X1+(g)→X2+(g)+e−X^{1+}(g) \rightarrow X^{2+}(g) + e^-

Factors Affecting Ionisation Energies

  1. Nuclear Charge: More protons, stronger attraction, harder to remove electrons.

  2. Distance from Nucleus: Electrons closer are harder to remove.

  3. Shielding Effect: Inner shells block outer electron attraction from nucleus.

  4. Ionic Charge: Higher charge on an atom increases difficulty in removing electrons.

Trends in Ionisation Energies

  • As you go down a group, ionisation energies decrease.

  • As you go across a period, ionisation energies increase, with jumps indicating the removal from a new inner shell.

General Trends

  • Group 2 elements: Be > Mg > Ca > Sr > Ba in 1st Ionisation Energy.

  • Period 3 elements: Na < Mg < Al < Si < P < S < Cl < Ar in 1st Ionisation Energy.