The Periodic Table Flashcards
What relationships are there between elements in the same:
- group?
- period?
- Groups: similar chemical properties
- Periods: trends in physical and chemical properties
Define first ionisation energy.
The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms..
What factors affect ionisation energy and why?
- Nuclear charge: the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons.
- Atomic radius: The closer the electron is to the nucleus, the stronger the nuclear attraction.
- Shielding: As the no. shells of electrons between the nucleus and valence electrons increases, the nuclear attraction on the valence electrons decreases.
What is the trend in 1st ionisation energy down a group and why?
- More shells of electrons so increased shielding effects.
- Atomic radius is larger so nuclear attraction is lower.
- These effects outweight the increase in nuclear charge, meaning that nuclear attraction on the valence electrons decreases overall.
- Less energy required to remove them therefore the 1st ionisation energies become less endothermic.
What is the general trend in first ionisation energies across a period?
- Proton number increases, increasing nuclear attraction.
- Atomic radius becomes smalller.
- Electrons are added to same shell so shielding doesn’t increase.
- Nuclear attraction therefore increases overall therefore the 1st ionisation energies become from endothermic across the period.
When going across a period, why is there a drop in first ionisation energies:
- from group 2 to 3?
- from group 5 to 6?
- Outer electron in Grp. 3 elements is in a p-orbital rather than an s-orbital, which has a slightly higher energy so the electron is slightly further from the nucleus. There is also additional shielding from the s electrons. These factors outweight the increased nuclear charge so the ionisation energy drops slightly.
- Outer electron in Grp. 6 is removed from a full orbital instead of a singly-occupied orbital, therefore it experiences spin-pair repulsion so the ionisation energy drops slightly.
Explain the trend in successive ionisation energies.
The energy per ionisation increases because:
- e- is removed from a shell closer to the nucleus/proton:electron ration is greater
- nuclear attraction is greater
- More E required to remove e-
How can you tell what group an element is in from its successive ionisation energies?
The number of IEs before a big jump (which signals the new electron is being removed from a shell closer to the nucleus) gives the number of valence electrons and therefore the group that the element is in.
What is a giant covalent lattice?
A network of atoms all bonded by strong covalent bonds.
What is are allotropes?
Different forms of the same element in the same state.
What is metallic bonding?
A strong electrostatic attraction between the metal cations and delocalised electrons.
What structure do solid metals form?
A giant lattice of closely packed cations surrounded by a sea of delocalised electrons.
What factors affect the:
- melting/boiling point
- malleability
- solubility
- thermal conductivity
- electrical conductivity
of metals?
- The no. of delocalised electrons per cation and the ionic radius.
- No bonds between cations therefore they are free to slide over each other, making metals malleable and ductile.
- No solid metals are soluble as the metallic bonds are too strong.
- The delocalised electrons can pass kinetic energy to each other, so metals are good thermal conductors.
- The delocalised electrons are free to move and carry charge, so metals are good electrical conductors.
What factors affect the:
- melting/boiling points
- thermal conductivity
- electrical conductivity
- solubility
of giant covalent substances?
- high melting point as lots of strong covalent bonds which all need to be broken
- vibrations travel easily through the stiff lattice so they are good thermal conductors
- all the electrons are held in localised bonds so they are not free to move and carry charge, so they are bad electrical conductors
- they are insoluble due to the strength of the covalent bonds
Explain the variation in melting and boiling points across a period.
- For the metals: the ionic radius decreases and ionic charge increases across the period so the metallic bonds get stronger so the m.p. and b.p. increase.
- For the giant covalent structures: high m.p. and b.p.
- For the simple molecular elements: m.p. and b.p. much lower as only weak IM forces need to be overcome. The more atoms per molecule that an element exists as (e.g. S8 vs. F2), the more dipersion forces there are and therefore the higher the m.p. and b.p.
