✨Module 3: Periodicity Flashcards
The more electrons in a molecule …
The stronger the London forces so higher b.m.p.
The higher the charge of an ion in an ionic compound …
The stronger the London forces so higher b.m.p.
Why does bromine have a low b.m.p?
Simple molecular with weak London forces BETWEEN molecules, which don’t require much energy to break.
Why does Ba have a low ionisation energy?
Increased distance/shielding outweighs increased nuclear charge.
Why does strontium metal have a high b.m.p?
It’s a giant metallic lattice with a regular arrangement of positive ions in a sea of delocalised electrons. Very strong forces between electrons and cations that require a lot of energy to break.
Define first ionisation energy.
Energy required to remove 1 electron (from each atom) in one mole of GASEOUS atoms to form one mole of GASEOUS 1+ ions.
3 factors that affect the ionisation energy.
Atomic radius - larger the distance between nucleus and outer electron, the lower the ionisation energy.
Nuclear charge - the more protons in the nucleus, the greater the attraction between nucleus and outer electron so bigger ionisation energy.
Electron shielding - inner shell electrons repel outer electrons so reduced ionisation energy as less nuclear attraction
Define second ionisation energy.
Energy required to remove 1 electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
Explain why the second ionisation energy of Helium is more than the first ionisation energy of helium.
After the first electron is lost, the single second electron is pulled closer to the nucleus as the number of protons don’t change (more + charge attracts the electron). Stronger nuclear attraction on remaining electron so more energy needed to remove it.
Mercury is the only metal that is a …
Liquid at room temp.
Define metallic bonding.
Strong electrostatic force of attraction between cations and delocalised electrons.
Cations are fixed in position, maintaining shape of metal.
Delocalised electrons can move throughout the structure.
Describe 3 factors that affect the strength of metallic bonding.
- No. protons - more protons mean stronger bonds.
- No. delocalised electrons per atom - more delocalised electrons mean stronger bonds. Transition metals generally have stronger bonds due to their d-electrons contributing to the electron sea
- Size of ion - the smaller the ion, stronger the bond. Smaller ions can pack more closely, increasing attraction between nuclei and delocalsed electrons.
- Higher charges on the metal ion attract the delocalised electrons more easily.
Explain why Mg has stronger metallic bonding that Na and therefore a higher b.m.p.
In Mg, there are more electrons in their outer shell that gets released to the sea of electrons. Mg ion is also smaller as it has one more proton. So stronger electrostatic attraction between positive metal ions and delocalised electrons and higher energy is needed to break the bonds.
Metals also …
Conduct electricity due to delocalised electrons that can move, and metals aren’t soluble.
Why don’t giant covalent substances conduct electricity?
Electrons can’t move. But in graphite, each carbon has one delocalised electron that can move.
Why are metals malleable?
Layers of positive ions can slide over each other.
Why is there a sharp decrease in melting point across period 2 + 3 between group 14 and group 15?
It marks a change from giant to simple molecular substances. Simple molecular substances like S8, P4, Cl2 have weak London forces between molecules, so little energy required to break them.
Ar is monatomic so weak London forces between ATOMS.
