ions and electrons Flashcards
Ionisation energy =
the amount of energy required to remove one mol of electrons from one mol of gaseous atoms of an element to form one mole of gaseous ions
unit of IE
kilojoules per mol
first IE
energy required to remove one mol of electrons from one mol of gaseous atoms of an element to form one mole of gaseous
1+ ions
calcium IE example
Ca (g) → Ca+ (g) + e-
Second IE =
energy required to remove one mol of electrons from one mol of gaseous 1+ ions to form one mol of gaseous 2+ ions
second IE example equation
X+ (g) → X2+ (g) + e-
OR (for calcium)
Ca+ (g) → Ca2+ (g) + e- IE2 = 1145 kJ mol-1
succesive ionisation energys increase
when you have removed the outer electron from an atom you have formed a positive ion
removing an electron from a positive ion is more difficult than from a neutral atom
as more electrons are removed attractive forces increase due to decreasing shielding and increasing proton to electron ratio
increase in IE is not constant
size of IE is affected by:
size of nuclear charge
distance of outer electrons from nucleus
shielding effect of inner electrons
spin-pair repulsion
ionisation energy across a period increases
nuclear charge increases meaning that the atomic radius of atoms decreases as the outer shell of electrons is pulled closer to nucleus due to a greater force of attraction
shielding remains constant as electrons are added to the same shell
it becomes harder to remove an electron as more energy is needed thus ionization energy increases
dips in the trend of ionisation energy across the period 2
decrease between beryllium and boron as the fith electron in born is in the 2p subshell which is further away from the nucleus than the 2s subshell of beryllium
decrease between nitrogen and oxygen due to spin pair repulsion - oxygen has 2 electrons in the 2p orbital so one of these elctrons pushes the other away making IE decrease as less energy is required to lose this electron
large decrease between periods of ionisation energy
increased distance from nucleus and outer shell as there is a whole new shell
imcreased shielding
these factors outweigh the increased nuclear charge
ionisation energy down a group decreases
number of protons increases so does nuclear charge
however
atomic radius increases as more shells are added thus giving more shielding so there is a weaker attraction meaning it is easier to lose e- so less IE needed (decreases)
if energy of welctrons is increased…
they jump to a higher energy level
this process id reversivble so electrons can return to their original energy levels - emitting energy when this occurs
absorbtion
into excited state, higher energy and further from the nucleus
emission
from excited state to ground state
line emmision spectra
each line is a specific energy value - suggesting that electrons can only possess a limited choice of allowed energies.
lines of energy are called quanta
lines gather at the blue higher energy end which shows the elctron is reaching its maximum energy
lines correspond to the electron jumping from higher energy levels down to the second energy level
succesive IE
first electron removed has a low IE as it is easily removed from the atom due to spin pair repulsion in the 4s orbital
second is more difficult as no spin pair repulsion
3rd is more difficult to remove as it is in a principal quantum shell that is closer to the nucleus
fourth is more difficult as the orbital is no longer full so less spin pair repulsion
large increase as electrons are removed from an increasingly positive ion
subshells contain …
orbitals
s orbital
spherical
size increases with increasing shell numver n=3 bigger than n=1
p orbital
dumbell shape
each shell has 3 p orbitals except the first shell
p orbitals can occupy x y and z axis
become larger and longer with increasing shell number
electrons have small spinning charges
creating a small magnetic field
electrons with the same spin repel each other
called spin pair repulsion
electrons into subshells (hunds rule)
must fill each orbital first
orbital can only have
2 electrons and they must have opposite spin
principal quantum shells increase in energy with
increasing quantum number eg. n=4 is higher in energy than n=2
orbitals in the subshell energy
is the same
electron configuration ghives information about:
number of electrons in each shell, subshell and orbital of an atom
principal quantum number> 1s^1 s=subshell small 1 =number of electrons
shorthand
includes symbol of nearest noble gas to account for how many e- are in that noble gas
hydrogen electronic config
1s^1
potassium shorthand
(Ar) 4s^1
chromium and copper electronic configurations (these are different than expected as they are energetically stable)
Using the Aufbau principle, you would write the following electron configurations
Cr = [Ar]
4s2 3d4
Cu = [Ar]
4s2
3d9
The actual electron configurations are:
Cr = [Ar]
4s1 3d5
Cu = [Ar]
4s1 3d10
beacuse electrons are lazy and want to have a full subshell. as they are more stable so therefore they both lose 1 e- in the 4s subshell and it goes to the 3d subshell
electrons are only paired when
there are no more empty orbitals available within a subshell
box notation
arrows must be pointing opposite directions
subshell location in periodic table
s - left
d- transition metals
p - right
f - weird ones at bottom
melting point trend in periods 2 and 3
from left to right
371, 923,932,1683,317,392,172,84
genral increase up untill silicon then significant decrease
due to the differences in bonding
across a period atomic ridius…
decreases
this is because the proton number and electron number increases by one every time you go an element to the right
same shielding
stronger attraction therefore pulled closer to the nucleus
why mp increases a bit avross a period
nuclear charge increases
atomic radius decreases
shielding is constant
harder to remove electron so more energy needed so IE increases
periodicity =
trend in repeating properties with increasing atomic number
