Periodic Trends Flashcards
What happens to ionisation energy across a period?
It increases because there is a greater effective nuclear charge for the same number of shells. There is also increased electrostatic attraction between valence electrons and the nucleus. So more difficult to remove an electron.
What happens to ionisation energy down a group?
It decreases because the valence electrons are further from the nucleus. There is greater shielding effect. Decreased electrostatic attraction between the valence electron and nucleus. It becomes easier to remove the electron.
What is electronegativity?
Ability of an atom to attract a pair of bonding electrons.
Factors that affect electronegativity (prioritise)
Number of shells, atomic radius (distance between nucleus and electrons involved in bonding), shielding, surface charge density, nuclear charge (smallest factor)
A has a lower electronegativity than B. What does this mean for A?
Indicates the attraction for the bonding pair in A in covalent compounds is less than the attraction for the bonding pair in B.
Why cations are smaller
Lost electrons from valence energy level and therefore has fewer shells. The distance between the nucleus and outer shell decreases. Both species have the same number of protons, but the cation has a greater effective nuclear charge. There is greater electrostatic attraction between the nucleus and valence electrons, pulling them in tighter. There is less repulsion between the remaining electrons (minor).
What happens to atomic radii across a period?
Decreases because there is an increase in nuclear charge (because more protons), whilst distance between the outermost energy level and the nucleus remains the same (because electrons are added to same energy level). Increased electrostatic attraction which draws the electrons in more tightly.
What happens to atomic radius down a group?
Increases because there are more electron shells surrounding the nucleus. Greatest factor is that valence electrons are much further from the nucleus, and are thus less attracted to the nucleus. Also increased repulsion between the electron shells (minor).
Why are anions larger than parent atoms?
More electrons in outer shell, and increased electron to electron repulsion which increases size of electron cloud.
What happens to electronegativity across a period?
Increases, because the nuclear charge increases whilst the number of electron shells doesn’t. Same shielding effect, so greater electrostatic attraction that attracts electrons more.
What happens to electronegativity down a group?
Decreases because the nuclear charge increases, but also increased number of electron shells which means distance between nucleus and outermost shell increases. Greater shielding effect, so less electrostatic attraction that attracts electrons less.
Why is ionisation energy endothermic?
Energy must be absorbed to overcome electrostatic attraction between nucleus and valence electron in order to remove it.
Define ionisation energy
Minimum amount of energy (in KJ) required to remove a single electron from an atom in the gaseous phase
Non-metals tend to have very high ionisation energies
Have more protons, and this increased nuclear attraction.
Why do the most reactive metals have low ionisation energies?
Metals react by losing electrons, so the more reactive a metal, the easier it is for it to lose an electron.
Account for the drop in ionisation energy from Be (4) to B(5) and Mg (12) to Al (13)
There is a decrease in ionisation energy despite an increase in the atomic number. B and Al have the outermost electron in a new p subshell. The p subshell is further from the nucleus than the s subshell. Because it is further away, there is less electrostatic attraction between the nucleus and valence electron, so less energy is required to remove it.
Why does Group 16 have lower ioinisation energy than Group 15?
Group 15 have 3 unpaired electrons in separate orbitals in a p subshell, which is a more stable arrangement. Group 16 have 2 unpaired electrons each in their own orbital, and 2 paired electrons in one full orbital. The paired electrons have slightly greater electron to electron repulsion, thus slightly less energy is required to remove an electron from a full orbital.
Large drop in ionisation from He to Li (at the end of the period)
Electron to be removed from He is in 1st energy level. Electron to be removed from Li is in second energy level. Although Li has more protons in nucleus (state) distance from the nucleus to outermost energy level is most influential factor determining electrostatic attraction of valence electrons, and thus ionisation energy.
Rise in ionisation energy from s1 to s2 (Cu to Zn and Na to Mg)
For both atoms, the electron removed will come off the same orbital (4s). However the atom with the s2 configuration has an extra proton in the nucleus which means there is greater electrostatic attraction between the nucleus and the valence electron which makes it more difficult to remove. There is a degree of repulsion between the electron pair in the s2 orbital, but not great enough to offset effect of the extra proton.
