Ionisation Flashcards
what is first ionisation energy
it is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms of an element to produce 1 mole of gaseous ions with a charge of +1
is ionisation endothermic or exothermic
endothermic because energy is needed to remove the electron
what is second ionisation energy
it is the energy needed to remove 1 mole of electrons from 1 mole of gaseous ions with a charge of +1 to produce 1 mole of gaseous ions with a charge of +2
why is there a general rise in the value of successive ionisation energies for an element
more energy is needed to remove an electron from a more positive ion
The proton:electron ratio increases each time an electron is removed so there is a greater effective nuclear charge
the ion becomes smaller so there is a greater attraction between the nucleus and the remaining electrons
what does a sharp rise on an ionisation energy graph show
an electron is removed from a new energy level which is closer to the nucleus
what does 2 sharp rises on an i.e. graph show
there are 3 energy levels
why is there a gradual increase in i.e on removal of electrons from the same shell
because the proton to electron ratio increases each time an electron is removed so there is a greater effective nuclear charge
why is the logarithm of the ionisation energy used in plotting an i.e. graph
the large range on numbers of ionisation energies is too big a range to plot directly, so taking logarithms makes the numbers manageable to plot
why is the first ionisation energy of selenium smaller than that of sulfur
the atomic radius of selenium is bigger than sulfur
there is more shielding of the outer electron in selenium
the nuclear attraction between the outer electron and the nucleus decreases so less energy is needed to remove the outer electron in selenium
how do you know which group an element belongs in based on its ionisation energies
if there is a large increase between the 3rd and 4th i.e it will belong in group 3
why does the atomic radius decrease across a period
the nuclear charge increases across a period because there are more protons
the shielding is the same across a period as the outer electrons are in the same shell
there is a greater nuclear attraction for the outer electrons across a period, so the electrons are held more tightly
why does the atomic radius increase as you go down a group
as you go down the group, the atoms have more electron shells, so the outer electrons are more shielded from the nuclear charge.
There are more protons as you go down the group but increased shielding outweighs this
there is less nuclear attraction for the outer electrons, so the electrons are held less tightly
why does the first ionisation energy increase across a period
the nuclear charge increases across a period because there are more protons.
The shielding is the same across a period, as the electron being removed is from the same shell
the atoms decrease in size across a period so there is a greater nuclear attraction for the outer electron and more energy is needed to remove it
why does the first ionisation energy decrease as you go down a group
as you go down the group, the atoms increase in size because they have more electron shells.
the outer electron is more shielded from the nuclear charge as you go down the group
There is less nuclear attraction for the outer electron so less energy is needed to remove it
why is there a sharp drop in first ionisation energy from one period to the next such as from Ar to K
the electron being removed from a potassium atom is from a new shell, n = 3, which is further from the nucleus and more shielded than the outer electron in an argon atom. Despite there being more protons in the nucleus of potassium, there is a less nuclear attraction for the outer electron in potassium
what are the 2 anomalies (groups) in ionisation energies
groups 2 + 3
AND
Groups 5 +6
what groups have the highest first ionisation energies
noble gases
what group has the lowest first ionisation energy
group 1
what happens to first ionisation energy across a period
it increase across a period
why is the first ionisation energy of chlorine higher than that of sodium
the nuclear charge is greater in a chlorine atom because it has more protons than sodium.
the shielding is the same in both atoms, as the electron lost in Na and Cl is in the same shell
the atomic radius of chlorine is smaller than sodium, so the nuclear attraction is stronger in a chlorine atom and more energy is needed to remove the outer electron in chlorine
why does helium have the highest first ionisation energy in the periodic table
helium has a greater nuclear charge than hydrogen because it has one more proton than hydrogen.
The shielding is same in both the atoms as the electron being removed in H and He is in the same sub-shell
Helium has a smaller atomic radius than hydrogen
as helium is the smallest atom in the periodic table, the electron being removed is closest to the nucleus so more energy is needed to remove it
how do you know about the ionisation energies in group 2 and 3 - the anomalies
write the electron configuration
how can tell the difference in ionisation energies in group 5 and 6 - the anomalies
draw the boxes and the arrows
why is the first ionisation energy of oxygen lower than that of nitrogen
the electron being removed from an oxygen atom is in a 2p orbital, containing 2 electrons.
There is repulsion between paired electrons, so less energy is needed to remove this electron
Further, nitrogen has a half filled sub shell which is more stable
why do group 0 element have the highest first ionisation energies in a period
they have the highest nuclear charge in a period.
the shielding is the same within a period
They have the smallest atomic radius in a period, so the nuclear attraction is strongest in these atoms
Further, these elements have a stable filled shell
why do group 1 elements have the lowest first ionisation energies in a period
they have the lowest nuclear charge in a period
the shielding is the same as the other elements in the period
they have the largest atomic radius in a period, so the nuclear attraction is weakest in these atoms, and less energy is needed to remove the outer electron
which group of element s would you expect to have the highest second ionisation energy
group 1 elements will have the highest 2nd i.e. as the second electron is removed from a new shell which is closer to the nucleus and shielded by fewer inner shells
isoelectronic
same number of electrons
why is the first i.e of Mg 2+ greater than Na + even though they are isoelectronic
Mg 2+ has a greater proton to electron ratio than sodium so magnesium has a greater effective nuclear charge than sodium
the shielding is the same in both the ions
Mg is smaller than Na so the outermost electron is more strongly attracted to the nucleus
