ions and electrons

Question 1

Ionisation energy =

Answer 1

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

Question 2

unit of IE

Answer 2

kilojoules per mol

Question 3

first IE

Answer 3

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

Question 4

calcium IE example

Answer 4

Ca (g) → Ca+ (g) + e-

Question 5

Second IE =

Answer 5

energy required to remove one mol of electrons from one mol of gaseous 1+ ions to form one mol of gaseous 2+ ions

Question 6

second IE example equation

Answer 6

X+ (g) → X2+ (g) + e-

OR (for calcium)

Ca+ (g) → Ca2+ (g) + e- IE2 = 1145 kJ mol-1

Question 7

succesive ionisation energys increase

Answer 7

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

Question 8

size of IE is affected by:

Answer 8

size of nuclear charge
distance of outer electrons from nucleus
shielding effect of inner electrons
spin-pair repulsion

Question 9

ionisation energy across a period increases

Answer 9

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

Question 10

dips in the trend of ionisation energy across the period 2

Answer 10

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

Question 11

large decrease between periods of ionisation energy

Answer 11

increased distance from nucleus and outer shell as there is a whole new shell

imcreased shielding

these factors outweigh the increased nuclear charge

Question 12

ionisation energy down a group decreases

Answer 12

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)

Question 13

if energy of welctrons is increased…

Answer 13

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

Question 14

absorbtion

Answer 14

into excited state, higher energy and further from the nucleus

Question 15

emission

Answer 15

from excited state to ground state

Question 16

line emmision spectra

Answer 16

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

Question 17

succesive IE

Answer 17

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

Question 18

subshells contain …

Answer 18

orbitals

Question 19

s orbital

Answer 19

spherical
size increases with increasing shell numver n=3 bigger than n=1

Question 20

p orbital

Answer 20

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

Question 21

electrons have small spinning charges

Answer 21

creating a small magnetic field

Question 22

electrons with the same spin repel each other

Answer 22

called spin pair repulsion

Question 23

electrons into subshells (hunds rule)

Answer 23

must fill each orbital first

Question 24

orbital can only have

Answer 24

2 electrons and they must have opposite spin

Question 25

principal quantum shells increase in energy with

Answer 25

increasing quantum number eg. n=4 is higher in energy than n=2

Question 26

orbitals in the subshell energy

Answer 26

is the same

Question 27

electron configuration ghives information about:

Answer 27

number of electrons in each shell, subshell and orbital of an atom principal quantum number> 1s^1 s=subshell small 1 =number of electrons

Question 28

shorthand

Answer 28

includes symbol of nearest noble gas to account for how many e- are in that noble gas

Question 29

hydrogen electronic config

Answer 29

1s^1

Question 30

potassium shorthand

Answer 30

(Ar) 4s^1

Question 31

chromium and copper electronic configurations (these are different than expected as they are energetically stable)

Answer 31

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

Question 32

electrons are only paired when

Answer 32

there are no more empty orbitals available within a subshell

Question 33

box notation

Answer 33

arrows must be pointing opposite directions

Question 34

subshell location in periodic table

Answer 34

s - left d- transition metals p - right f - weird ones at bottom

Question 35

melting point trend in periods 2 and 3

Answer 35

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

Question 36

across a period atomic ridius...

Answer 36

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

Question 37

why mp increases a bit avross a period

Answer 37

nuclear charge increases atomic radius decreases shielding is constant harder to remove electron so more energy needed so IE increases

Question 38

periodicity =

Answer 38

trend in repeating properties with increasing atomic number