Periodicity Flashcards
Why does oxygen have a lower first ionisation energy than nitrogen?
Oxygen is 1s2 2s2 2p4 & nitrogen is 1s2 2s2 2p 3.
With the box arrows thing oxygen fills up /\ /\ /\ / /. (Updown updown updown up up) and nitrogen fills up /\ /\ / / / (updown updown up up up)
As you move from nitrogen to oxygen you have 4 electrons in the p sub shell instead of 3. This means that there will be one pair because there are only three p orbitals in a p sub shell. These two electrons repel each other, increasing the atomic radius and decreasing the electrostatic force of attraction between the nucleus and the electron. Therefore this decreases the amount of energy need to overcome this attraction.
Why is the first ionisation energy of boron less than the ionisation energy of beryllium?
Beryllium is 1s2 2s2 & boron is 1s2 2s2 2p1
With the box arrows thing beryllium fills up /\ /\ (updown updown) and boron fills up as /\ /\ / (updown updown up)
As you move from beryllium to boron there is a 5th electron in the 2p subshell. This means the electron configuration of boron is less stable than beryllium’s therefore less energy is required to remove one electron from boron. Beryllium has a stable configuration with the 2s sub shell completely full so more energy is required to remove an electron.
Why does sulphur have a lower first ionisation energy than phosphorus?
As you move from phosphorus to sulphur you have 16 electrons in the 3p subshell instead of 15. The 16th electron fills up 3p subshell and it means the will be one pair of electrons because there are only 3 p orbitals in a p subshell. These two electrons repel each other, increasing the atomic radius and decreasing the electrostatic force of attraction between the nucleus and the electron. Therefore this decreases the amount of energy need to overcome this attraction.
Why is the first ionisation energy of gallium significantly lower than that of zinc?
Because the last electron (the 31st) of gallium is in the 4p orbital instead of, like zinc where the last electron (the 30th) is in the 4s orbital. When you move from the 4s orbital out to the 4p orbital you move slightly further out from the nucleus AND the 4p orbital is shielded by the 4s orbital so this means the distance from the nucleus (atomic radius) increases, so that means the electrostatic attraction between the nucleus and electrons is weaker which means that less energy is needed to overcome that attraction and remove the electron.
Why does lithium have a lower first ionisation energy than neon?
Lithium, although it has the same number of shells and similar shielding it has a lower atomic number than neon which means lower nuclear charge than neon. This means that there is a weaker attraction between the nucleus and outer shell electrons and therefore it requires less energy to overcome this.
Explain the general trend in ionisation energy across period 2?
Across the period there is the same number of shells so SIMILAR SHIELDING(remember to say this!!!). As nuclear charges increases it causes a stronger force of attraction on the outer most electrons so more energy is required to overcome the electrostatic attraction.
Explain the general trend in ionisation energy across period 3?
Same as period 2 - Across the period there is the same number of shells so similar shielding. As nuclear charges increases it causes a stronger force of attraction on the outer most electrons so more energy is required to overcome the electrostatic attraction.
What areas of the periodic table have trends and acceptions in ionisation energies?
General trend across period 2 (acceptions Be - B)
General trend across period 3 (acceptions Mg - Al)
Mg to Al
Aluminium’s outer electron is in a 3 p orbital rather than 3s like in Mg, so the electrons are found slightly further from the nucleus and the 3p orbital has additional shielding provided by the 3s2 electrons, these factors (shielding and distance) are strong enough to override effect of increased nuclear charge so we get a drop in ionisation energy
Neon to helium
Down group new shell of electrons so further out, increases shielding, larger atomic radius, weaker attraction, lower IE
Going across the periods 2 and 3 from left to right
Increase in melting and boiling points because the atomic number increases whereas the shielding remains the same so attraction between the two nuclei gets stronger.
Going down the group melting and boiling point????????
Decreases because extra shell of electrons and larger atomic radius due to more shielding so greater distance between the nuclei of the atoms and so weaker force of electrostatic attraction so less energy required to overcome these forces.
If you are comparing melting points what should you think about?
Shielding (shielding and electrons)
Atomic radius
Nuclear charge
[Ionisation graph] Going down a long way
Eg. He down to Li
New group - additional shell - more shielding and electrons - larger atomic radii- weaker attraction between nucleus and outer shell electrons- less ionisation energy required to over come this attraction
[Ionisation graph] Going up a long way
Eg. Na up to Ar
Going across groups - increasing atomic number - increasing nuclear charge - same number of shells - stronger attraction between nucleus and outer shell electrons - higher ionisation energy required to break this attraction.
[Ionisation graph] Which jumps downwards are due to repelling?
P down to S
N down to O
As down to Se
[Ionisation graph] Which jumps downwards are down to a further out p orbital and shielding of electrons?
Mg to Al
Zn to Ga
Why does Selenium have a lower ionisation energy than Arsenic?
Selenium is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 and has 34 electrons /\ /\ /\/\/\ /\ /\/\/\ /\/\/\/\/\ /\ /\ / /
Arsenic is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 and has 33 electrons /\ /\ /\/\/\ /\ /\/\/\ /\/\/\/\/\ /\ / / /
The 34th electron of selenium has to go into the 4p orbital because it can hold six electrons. This means there is one pair of electrons in the 4p orbital compared to arsenic which doesn’t have any pairs of electrons as only 3 electrons are in its 4p orbital. These two electrons repel each other, increasing the atomic radius and decreasing the electrostatic force of attraction between the nucleus and the electron. Therefore this decreases the amount of energy need to overcome this attraction.
[ionisation graph] Which jumps downwards are due to unstable electron configurations?
Be down to B
Ionisation energy is always
EXOTHERMIC
Why do Li and Na have similar ionisation energies?
There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater in Sodium. However, sodium doesn’t have a much larger IE because this increased nuclear charge is cancelled out by sodium’s outer shell’s electron greater distance from the nucleus + additional shielding of shells and electrons which means a weaker electrostatic attraction between the e- and nucleus.
Why is there a large jump between the sixth and seventh ionisation energy of oxygen?
Oxygen has 8 electrons,( /\ ) - 1s part of first circle (/\ /\ / /) - 2s and 2p part of second circle. The first six removed starting from the outer shell are more easily removed because they are being removed from the same shell. However the 7th electron moves into a different shell, which is the closest one to the nucleus. This means there is a smaller ionic radius and so there is a very strong electrostatic attraction between this electron and the nucleus. There are also no more electrons shielding the 7th electron. Therefore lots more energy is required to overcome this attraction.
First IE equation
X(g) —> X+(g) + e-
2nd IE equation
X+(g) —> X2+(g) + e-
