F321 - Topic 1 Atoms and Electron Structure

Atomic Structure

• The nucleus contains almost all of the mass of an atom, as it cotains protons and neutrons.
• The nucleus of an atom contains all of the positive charge.
• The electrons orbit outside the nucleus.
• Electrons are negatively charged.

Definitions

• Atomic number = number of protons in the nucleus.
• Mass number = number of protons and neutrons in the nucleus.
• Number of neutrons = mass number - atomic number.
• Number of electrons = number of protons (in a neutral atom)

Isotopes

• Isotopes are atoms with the same number of protons but different number of neutrons.
• Isotopes of an element have the same chemical properties, as they have the same electron arangement.

Ions

• Ions are formed when atoms gain or lose elctrons.
• E.g. An atom Cl has 17 electrons, a Cl- ion has gained one electron.

Relative Atomic Mass

• Average mass of an atom relative to 1/12 of the mass of a carbon-12 atom.
• To calculate relative atomic mass, add together (mass number x percentage/100).

Example

• 75% of Cl atoms have a mass number of 35
• 25% of Cl atoms have a mass number of 37

• Average mass of a Cl atom    = (mass no x percent/100) + (mass no x percent/100)

                                                      = (35 x 75/100) + (37 x 25/100)

                                                      = 35.5

Relative Isotopic Mass

• Mass of a particular isotope relative to 1/12 of the mass of a carbon-12 atom.

Orbitals

• An orbital is a region that can hold up to two electrons with opposite spins.
• Orbitals have different shapes called - s, p, d and f.

S orbitals

•  S orbitals are spherical in shape and come in sets of one.

P orbitals

• P orbitals are hour - glass shaped and come in stes of three.

D orbitals

• D orbitals come in sets of five which can hold up to 10 electrons.

E.g.    Iron - 26 electrons - 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Successive Ionisation Energies

• Energy change when one mole of electrons is removed from one mole of gaseous positively charged ions.
• In an element, successive ionisation energies get bigger because the emaining electrons are held more tightly by the unchanged nuclear charge.

First Ionisation Energy

• Energy change when one mole of electrons is removed from one mole of gaseous atoms.
• M(g) -----> M⁺(g) + e^-
• Ionisation is easier if :
  • The nuclear charge is smaller.
  • The electron is further away from the nucleus.
  • There is more shielding.

• First ionisation energy shows periodicity.

Factors affecting the size of ionisation energy

• The size of that attraction will be governed by:
  • The charge on the nucleus.
  • The distance of the electron from the nucleus.
  • The number of electrons between the outer electrons and the nucleus (shielding).
  • Whether the electron is on its own in an orbital ( paired electrons are easier to remove).

• Why the ionisation energy increases:
  • There is an increasing number of protons in the nuceus.
  • The increasing nuclear charge, drags the outer electrons in closer to the nucleus.

• Why:
  • There is a bigger number of protons.
  • But, the shielding and nuclear radius overcomes the nuclear charge.
• E.g.
• • Li -1s²2s¹1st I.E. = 519 kJ mol^-1
  • Na -1s²2s²2p⁶3s¹1st I.E. = 494 kJ mol^-1

Blocks

• The s block is group 1 and 2.
• The p block is group 3 and 8.
• The d block is between the s and p blocks.