Periodicity Flashcards
(15 cards)
State and explain the trend in atomic radius across a period
Atomic radius decreases due similar shielding present, yet an increase in the electrostatic force of attraction between nucleus and outer electron shell due to an increase in nuclear charge since proton number is increasing
Describe the trend in melting point across period 3
Between Na and Al:
General increase as metal ions have an increasing positive charge, increasing number of delocalised electrons and smaller atomic radius. This means an increasing strength of the metallic bond’
Si:
Highest M.P as it is giant covalent, lots of energy needed to overcome the many strong covalent bonds between atoms
P:
Much lower M.P than Si as it is simple molecular. Little energy is required to overcome the weak VDW forces between molecules
S to Ar
S has a slightly higher MP than P as it is a slightly bigger molecule, therefore stronger VDW forces
Trend from S to Ar is a decrease due to weaker VDW forces (Cl2 is a smaller molecule than P4 or S8, and Ar only exists as individual atoms)
Why does Ar have the lowest melting point across period 3
Exists as a single atom Ar which has less e- than S8 or P4 or Cl2, which are also simple molecular but have stronger VDW
What is ionisation energy
The enthalpy change required to remove one mole of electrons from one mole of gaseous atoms
What 3 factors affect ionisation energy
Shielding (increased shielding means lower ionisation energy)
Nuclear charge (Increase in nuclear charge means higher ionisation energy
Atomic radius (Increase in atomic radius means a lower ionisation energy
What is periodicity
The repeating pattern of physical or
chemical properties going across the periods
What is the trend in successive ionisation energies, and where do we get the biggest jumps
Trend is an increase in ionisation energy successively due to electrons being removed from a more positive ion when comparing electrons being removed from the same orbital
When comparing electrons from different orbitals, this is where we get the biggest jumps as removing electrons from an orbital closer to the nucleus requires much more energy due to an increase in the electrostatic force of attraction between positive nucleus and outer electron
Explain the trend in ionisation energy down a group
Ionisation energy generally decreases
Increase in atomic radius
Increase in shielding
Weaker electrostatic force of attraction between nucleus and outer electron
What does decreasing ionisation energy down a group prove
Proves that electrons are in shells
State and explain the general trend in ionisation energy across period 3, what are the exceptions
General increase:
Increase in the number of protons, shielding is similar and atomic radius marginally decreases therefore more energy is required to remove an outer electron
Al and S are the exceptions
Explain the exceptions to the trend in ionisation energies across period 3
Al - The outer electron in Al sits in a higher energy 3p subshell which is slightly further from the nucleus and therefore has a lower first ionisation energy compared to Mg, whose outer electron sits in the 3s subshell
S - The outer electron in S sits in an orbital with 2 electrons in it. Thus, as opposite charges repel, less energy is required to remove an electron from an orbital with 2 electrons compared to an orbital with 1 electron such as in P
What affects ionisation energy the most
What affects the strength of an ionic bond the most
Ionisation energy is more affected by atomic radius, since a larger distance weakens the attraction between the nucleus and the outer electron.
Ionic bonding is more affected by ionic charge, as a higher charge increases the strength of electrostatic attraction between ions.
How does atomic radius affect bond enthalpy?
A larger atomic radius means the bonding electrons are further from the nucleus, resulting in weaker attraction and lower bond enthalpy. Smaller atoms form stronger bonds because their nuclei attract the shared electrons more strongly.
What factors affect the strength of an ionic bond, in order of priority
Ionic Charge
Ionic Radius
Charge Density
What can we think of ionic attraction as
Centre of cation to centre of anion
(Hence why increasing atomic radius decreases ionic attraction)
In reality, ionic attraction arises from the electrostatic force between entire charge clouds, not just nucleus-to-nucleus but for A-Level this is fine
