Periodicity Flashcards
How is the periodic table organised
In increasing atomic number
Give the definition of ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous atoms (under standard conditions).
Give the general first ionisation energy equation of a metal
M (g) –> M^+(g) + e^-
Give the general second ionisation energy equation of a metal
M^+ (g) –> M^2+(g) + e^-
What happens to the ionisation energy of group 1 alkali metals as you go down the group and why
Reactivity increases down the group because as you go down the group the atoms get bigger and so there is a greater distance between he outermost shell and the nucleus
Is the 2nd ionisation energy higher or lower than the 1st - why
Higher because the overall charge on the atom is more positive. There is greater attraction and hold over the electron being removed than before, requiring more energy to overcome.
Across a period, what is the general trend in ionisation energy? Give two reasons
Increases
-Increase in nuclear charge along group
-BUT same shell (distance from nucleus) so shielding remains fairly constant
How can you identify an element from a graph of its successive ionisation energies?
How many ionisation energies there are before the initial jump gives you the group of the element. The number of levels give you the period that the element is in.
What are the four ‘blocks’ of the periodic table, and how many electrons do each hold?
S - 2
P - 6
D - 10
F - 14
What are the names of the groups below the periodic table?
Lanthanides
Actinides
What elements are metalloids (list/circle on periodic table)
B, SI, G, As, Sb, Te, Po
Is Zinc a transition metal?
No - it has a full d-shell
Why is the Be 1st I.E more than the B?
Electron configuration
the fifth electron in boron is in the 2p subshell, which is further away from the nucleus than the 2s subshell of beryllium.
Why is the N 1st I.E more than the O?
The paired electrons in the 2p subshell of oxygen repel each other, making it easier to remove an electron in oxygen than nitrogen.
Ionisation energy trend down a group (three reasons)
Decreases
-More shielding between outermost electron being removed and positive nucleus
-Bigger distance due to more shells
What is electron affinity?
The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous ions.
General trend of electron affinity down the group
Decreases
As you go down the group there are more shells and shielding and so it is harder for the electron to be attracted by the nucleus
General trend of atomic radius across the period
Decreases because nuclear charge increases, pulling the electron shells closer
General trend of atomic radius down the group
Increases due to the increasing number of shells
Ionic radii compared to atomic radii in atoms that form positive ions
smaller (electrons lost)
Ionic radii compared to atomic radii in atoms that form negative ions
bigger (electrons gained)
