Atomic Structure

relative atomic mass definition

the average mass of an atom of an element compared to 1/12th of the mass of a Carbon-12 isotope

properties of isotopes

chemical properties are identical as they have the same number and arrangement of electrons
physical properties may differ due to different masses

relative abundance formula

Ar = (mass x abundance)... / total abundance

finding the % abundance

s sub-shells

contain 1 s orbital holding a max of 2 electrons

p sub-shells

contain 3 p orbitals holding a max of 6 electrons

d sub-shells

contain 5 d orbitals holding a max of 10 electrons

f sub-shells

contain 7 f orbitals holding a max of 14 electrons

filling up energy levels

electrons occupy the lowest available energy level first (Aufbau principle)

electronic configuration and the periodic table

an element's location on the periodic table gives its complete electronic arrangement - look on the table to find its outermost sub-shell (however far it is into the section is how many electrons in that shell)
all sub-shells before then are full

box and arrow notation

each box represents one orbital, each arrow represents an electron
always label the boxes
electrons prefer to be in an orbital alone, fill up all orbitals in subshell with one electron before pairing them

first ionisation energy

energy required to remove an electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions

first ionisation energy equation

X (g) -> X+(g) + e-

always include state symbols

second ionisation energy

energy required to remove an electron from each ion in 1 mole of gaseous +1 ions to form 1 mole of gaseous 2+ ions

second ionisation energy equation

X+ (g) -> X 2+ (g) + e-

is ionisation an endothermic or exothermic process?

removing electrons from atoms is an endothermic process as energy is needed to overcome the attraction of the nucleus on the outgoing electron. The stronger the attraction, the greater the ionisation energy

3 factors that affect ionisation energy

1) distance of the electron from the nucleus (closer to nucleus means stronger attraction so higher ionisation energy)

2) number of inner electron shells (repel outer electrons and shield positive charge of nucleus meaning lower ionisation energy)

3) number of protons in nucleus (more protons means greater attraction)

successive ionisation energies

how much energy is required to remove an electron from an atom until all electrons are removed

why are log ionisation energies plotted on successive ionisation energy graphs?

there are large differences between successive values (large range)

log=to power of 10

during successive ionisation, why do electrons in different shells have large differences in energy required?

as they are in different shells, they are different distances from the nucleus. Electrons in higher shells experience less attraction as they are further from the positive nucleus and there are more inner shells shielding them. This means that less energy is required to remove them

during successive ionisation, why do electrons in the same principle energy shell have similar values for energy but with small incremental increases?

as they are all in the same shell, they are the same distance from the nucleus with the same amount of inner level shielding. However, as each electron is removed, the ion becomes increasingly positive, meaning more energy is required to overcome the attraction (explaining the small differences)

trends in first ionisation down the groups in the periodic table

as you go down any group, 1st ionisation energy gets lower

although the number of protons in the nucleus increases, the electron being removed is further from the nucleus and experiences more inner electron shielding so less energy is required

patterns in atomic radius across a period

across a period, all atoms have the same number of shells and experience the same amount of shielding but the nuclear charge increases as there are more protons. The electrons are more strongly attracted to the nucleus, so the atomic radius decreases

which elements form ions which are smaller than the atom from which they are formed?

elements that form cations as they lose the outer shell

which elements form ions which are larger than the atom from which they are formed?

elements that form anions have a bigger atomic radius

they don’t gain shells (they fill the outer shell) yet the electrons repel from each other making the ion larger

isoelectronic

same electronic configuration

trends in first ionisation energy across a period

general increase as proton number increases (more attraction)

there are drops after new shells and sub shells with significant drops between periods

explain the drop between magnesium and aluminium

from an s -> p sub shell
p sub shells are further from nucleus so less energy is needed to remove electrons

explain the drop between phosphorous and sulphur

this is the 1st time electrons are paired in p sub shell

increased repulsion so less energy is needed to remove electrons

in period 2, why is the first ionisation energy of oxygen less than that of nitrogen?

in oxygen, electron is removed from a paired 2p orbital whereas in nitrogen, the electron is removed from a singularly occupied orbital

X and Zn are different elements. Explain why the chemical properties of 70-X and 70-Zn

they don’t have the same electronic configuration