3.1.1 Periodicity Flashcards
Define periodicity
The trends in elements properties with increasing atomic number.
How is the periodic table arranged?
- by increasing atomic number
- in periods with repeating trends in physical and chemical properties
- in groups having similar chemical properties
Examples chemical properties
toxicity, flammability, oxidation states, reactivity, chemical stability etc…
Examples of physical properties
colour, solubility, odour, hardness, density, melting point and boiling point.
Define first ionisation energy
The amount of energy needed to remove 1 mole of electrons from gaseous atoms (to form 1 mole of 1+ gaseous atoms).
State the 3 factors that affect ionisation energy
- Atomic radius
- nuclear charge
- shielding
What is ionisation energy the measure of?
-how easily an atom loses electrons to form positive ions
What is atomic radius?
The distance between the nucleus and the outer-shell electrons.
What is nuclear charge?
The attraction between the protons in the nucleus and the outer-shell electrons.
What is electron shielding?
The repulsion between inner-shell electrons and outer-shell electrons. More shells = more shielding.
Describe the trend in ionisation energy across a period from left to right.
Ionisation energy increases as you go across a period:
- nuclear charge increases as there are less electrons, so the atoms are smaller
- atomic radius decreases due to a decrease in shells
- shielding is the same
- nuclear attraction is increased
Describe the trend in ionisation energy down a group.
Ionisation energy decreases as you go down a group:
- nuclear charge decreases as there are more electrons
- atomic radius increases
- increased shielding
- nuclear attraction decreases
What is the equation for first ionisation energy?
M(g)—>M+(g) + e-
What is the equation for second ionisation energy?
M+(g)—> M2+(g) + 2e-
Which groups are in S block?
Groups 1-2
Which groups are in D block?
Groups 3-12
Which groups are in the P block?
Groups 13-18
Describe the trend in ionisation energy across period 1.
- Hydrogen: despite having a nuclear charge of 1+, it has a relatively high 1st ionisation energy as its electron is closest to the nucleus and has no shielding.
- Helium: a much higher value because of the extra proton in the nucleus. Additional proton = stronger attraction= harder to remove electron.
Describe the trend in ionisation energy across period 2.
- Lithium: from helium there is a substantial drop because there is an extra electron in the next energy level. Despite the increased nuclear charge the effective nuclear charge is less because of the shielding effect of the filled inner 1s energy level. 2s is also further from the nucleus.
- Beryllium: higher ionisation energy than lithium due to increased nuclear charge.
Continued trend in ionisation energy across period 2
Boron
- Boron: ionisation energy drops from beryllium, this is because of the extra electron in one of the 2p orbitals. Increased shielding makes it easier to remove the electron.
- Carbon: ionisation energy increases again due to increased nuclear charge. The extra electron does not pair up in the same orbital but occupies another 2p orbital.
Continued trend in ionisation energies across period 2
Nitrogen
- increase in ionisation energy from carbon due to increased nuclear charge. Extra electron goes into a different 2p orbital (3 unpaired electrons in 2p orbitals)
- oxygen: there is a drop. The extra electron pair, pairs up with the first 2p orbital.
Continued trend in ionisation energy across period 2 (fluorine)
- Fluorine: ionisation energy increases due to increased nuclear charge
- Neon: ionisation energy increases due to increased nuclear charge. 2p orbitals are now full
