The Periodic Table

The how and whyThe how and why

2

Atomic Size First problem where do you start

measuring. The electron cloud doesn’t have a

definite edge. Chemists get around this by measuring

more than 1 atom at a time.

3

Atomic Size

Atomic Radius = half the distance between two nuclei of a diatomic molecule (homo-nuclear molecule).

}Radius

4

Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level is further away. The effective charge from the

nucleus The greater the nuclear charge

reaching the valence electrons the closer these electrons are pulled in.

5

Group trends As we go down a group Each atom has another

energy level, so the atoms get bigger.

There are more levels in the kernel and therefore greater shielding of valence electrons (weaker attraction).

HLi

Na

K

Rb

6

Periodic Trends As you go across a period the radius gets smaller. Valence electrons are in the same energy level. Greater nuclear charge gets to the valence electrons –

same shielding e- but greater nuclear charge. Outermost electrons are closer. Row 3 elements all have 10 shielding electrons; +11 (nuclear charges) +18

Na Mg Al Si P S Cl Ar

7

Overall

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

8

Ionization Energy The amount of energy required to

completely remove an electron from a gaseous atom.

Removing one electron makes a +1 ion.

The energy required is called the first ionization energy.

9

Ionization Energy The second ionization energy is the

energy required to remove the second electron.

Always greater than first IE. The third IE is the energy required to

remove a third electron. Greater than 1st of 2nd IE.

10

Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11749 14840 3569 4619 4577 5301 6045 6276

11

Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11749 14840 3569 4619 4577 5301 6045 6276

12

What determines IE The greater the effective nuclear

charge the greater IE. Less distance from nucleus

increases IE Filled and half filled sublevels have

lower energy, so removing them raises the IE.

Shielding: effective blocking of nuclear charge weakens the attraction of valence electrons.

13

Shielding The electron on the

outside energy level has to look through all the other energy levels to see the nucleus

14

Shielding The electron on the

outside energy level has to look through all the other energy levels to see the nucleus.

A second electron has the same shielding.

15

Group trends As you go down a group first IE

decreases because The electron is further away. More effective shielding.

16

Periodic trends All the atoms in the same period have

the same energy level. Same shielding. Increasing nuclear charge So IE generally increases from left to

right. Exceptions at full and 1/2 fill sublevels.

17

Firs

t Ion

izat

ion

ener

gy

Atomic number

He

He has a greater IE than H. no shielding greater nuclear charge

H

18

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He Li has lower IE than H more shielding further away (outweighs greater

nuclear charge)

Li

19

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE than Li same shielding greater nuclear charge removing an electron

from a full s sublevel

Li

Be

20

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He B has lower IE than Be same shielding greater nuclear charge removing an electron from a

partially filled p sublevel

Li

Be

B

21

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

22

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

23

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O Breaks the pattern

because in N the electron gets removed from a ½ filled p sublevel

24

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

25

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Ne has a lower IE

than He Both have full

valence shells Ne has more

shielding Greater distance

26

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Na has a lower

IE than Li Both are s1

Na has more shielding

Greater distance

Na

27

Firs

t Ion

izat

ion

ener

gy

Atomic number

28

Driving Force Full Energy Levels are very low

energy. Noble Gases have full energy levels. Atoms behave in ways to achieve

noble gas configuration (become isoelectronic).

29

2nd Ionization Energy For elements that reach a filled or

half filled sublevel by removing 2 electrons 2nd IE is lower than expected.

True for s2

Alkaline earth metals form +2 ions.

30

3rd IE Using the same logic s2p1

atoms have an low 3rd IE.

Atoms in the aluminum family form +3 ions.

2nd IE and 3rd IE are always higher than 1st IE!!!

31

Electron Affinity The energy change associated with

adding an electron to a gaseous atom. Easiest to add to group 17 (7A). Gets them to full energy level becomes

stable – releases a large amount of energy.

Increase from left to right atoms –metals are losers so to gain an electron will only bring about a small increase in stability (if any) – small energy change.

Decrease as we go down a group.

32

Ionic Size Cations form by losing electrons. Cations are smaller than the atom

they come from. Metals form cations. Cations of representative elements

have noble gas configuration.

33

Ionic size Anions form by gaining electrons. Anions are bigger than the atom they

come from. Nonmetals form anions. Anions of representative elements

have noble gas configuration.

34

Configuration of Ions Ions always have noble gas

configuration. Na is 1s22s22p63s1 Forms a +1 ion - 1s22s22p6 Same configuration as neon. Metals form ions with the

configuration of the noble gas before them - they lose electrons.

35

Configuration of Ions Non-metals form ions by gaining

electrons to achieve noble gas configuration.

They end up with the configuration of the noble gas after them.

36

Group trends Adding energy level Ions get bigger as

you go down.

Li+1

Na+1

K+1

Rb+1

Cs+1

37

Periodic Trends Across the period nuclear charge

increases so they get smaller. Energy level changes between

anions and cations.

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

38

Radii vs atomic #

3

45 6

78 9

11

12

1314

1516

17

19

20

31 32 33 34 35

37

38

4950 51 52

3

4

56

7

8 9

11

12

1314

15

16 17

19

20

3132

33

34 35

37

38

4950

51

52

0.01

0.06

0.11

0.16

0.21

3 4 5 6 7 8 9 11 12 13 14 15 16 17 19 20 31 32 33 34 35 37 38 49 50 51 52

Atomic #

Rad

ii (n

onom

eter

s)atomic radii ionic radii

39

Size of Isoelectronic ions Iso - same Iso electronic ions have the same #

of electrons (same configuration) Al+3 Mg+2 Na+1 (Ne) F-1 O-2 and N-3 all have 10 electrons all have the configuration 1s12s22p6

40

Size of Isoelectronic ions Positive ions have more protons

than electrons so they are smaller (strong attraction for electrons).

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

Electronegativity

42

Electronegativity The tendency for an atom to attract

electrons to itself when it is chemically combined with another element.

How fair it shares. Big electronegativity means it pulls

the electron toward it. Atoms with large negative electron

affinity have larger electronegativity.

43

Group Trend The further down a group the farther

the electron is away and the more electrons an atom has – the lower the electronegativity value.

44

Periodic Trend Metals are at the left end. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away. High electronegativity.

45

Ionization energy, electronegativity

Electron affinity INCREASE

46

Atomic size decreases, shielding constant

Ionic size decreases

Atomic size increases

Shielding increases