p-block - periodicity + physical properties

elements which have their valence electrons in p-orbitals

non-metals, metalloids, metals

atomic radii tends to decrease across a period due to increasing effective nuclear force

atomic radii tends to increase down a group because of increasing principle QN meaning more shells, however increase down group is not linear, the change from row 2 to row 3 is significantly larger, this tells us row 2 elements are significantly smaller than row 3

based on bonding atoms of that element, so elements that aren’t known to bond will not have these values

the ability of an element to attract electrons in a chemical bond

electronegativity tends to increase across a period as effective nuclear charge increases and atomic radii decreases, leading to an increase in the attraction between nucleus and bonding electrons

electronegativity tends to decrease down a group as bonding electrons are further from the nucleus as atoms have more electron shells

they don’t form bonds

metals X <~2.00 metalloid X ~ 2.00 ~ 2.20 non metal X > ~ 2.20 - this is relatively accurate

the energy required to homolytically break a single bond between 2 atoms of the same element, represents bond strength - as homolytic fission this creates 2 radical species

tends to increase across a period as electronegativity increases across a period, so elements attract bonding electrons more strongly so bonds get stronger - exceptionally group 14 –>15 sees a decrease in bond energy because group 15 elements still have lone pairs when saturated with bonds, which repel and weaken the bond

tends to decrease down a group, except for between N->P, O->S, F->Cl this is because p-orbitals have n-2 radial nodes, so down a group radial nodes increases as principle QN n increases, reducing in-phase orbital overlap when bonds form, resulting in weaker bonds - exceptions is because lone pairs on N, O and F weakens their homonuclear single bonds, and as they also have very small atomic radii they have very short bond lengths so lone pair repulsion is greater, compared to row 3

most p-block elements do not form double bonds

group 13 elements have only 3 valence electrons, typically held in sp2 orbitals leaving an empty p-orbital which cannot form a π-bond as its empty

group 17+18 don’t need to form multiple bonds, due to the octet rule, both achieve an octet either naturally or after just one single bond

less likely as row 2 elements have very small atomic radii their single bonds are shorter, allowing atoms to be close enough to allow effective π-orbital overlap elements from row 3 down become much larger so π-orbital overlap is poorer - some row 3 elements can be forced to form double/triple bonds

P + S are the smallest row 3 elements able to form multiple bonds, groups 13 + 17 + 18 aren’t able to at all and the atomic radii of Si is too large so overlap is much poorer, although it ca sometimes be forced P + S aren’t yet that big and so have decent π-orbital overlap

yes

this is based on the energy difference for forming a single bond vs multiple bonds - for N and O forming triple/doble bonds is favourable as their single bonds are much weaker due to lone pair repulsion, and they have good π-orbital overlap as they have short bond lengths, the energy cost of a double bond is < energy cost of 2 single bonds - same for triple - for P and S, π-orbital overlap isn’t as good and their single bonds are stronger, the energy cost of double bond > energy cost of 2 single bonds, so they prefer to form single bonds

O=O N≡N P-P S-S (remember P and S don’t necessarily exist as diatomics)

in the absence of strong intermolecular interactions, m.p/b.p generally increases with Mr compounds with ionic/polymeric structures have high m.p/b.p that compounds with discrete molecular structures compounds that can hydrogen bond also have much higher m.p/b.p

proton donor

proton acceptor