Chemical Periodicity

OBJECTIVE:• Interpret group trends in atomic

radii, ionic radii, ionization energies, and electronegativities.

Periodic TrendsPeriodic Trends

What nuclear charge is felt by the outermost electrons of a magnesium atom?

1s2 2s2 2p6 3s2

Slater’s Rules

How much energy is needed to remove thelast electron from a magnesium atom?

1s2 2s2 2p6 3s2 → 1s2 2s2 2p6 3s1 + e-

22

422n

hnemZ2E π

−=

Slater’s Rules

1s2 2s2 2p6 3s2

blocks 85% of outside3s electron

blocks 100%of outside 3s

electron

Z* = Zeff = Z - S

= 12 - [2(1.00) + 8 (0.85) + 1 (0.35)] = 2.85

blocks 35% of theother 3s electron

Slater’s Rules

22

422n

hnemZ2E π

−=

2

21

2

2H

nZ)molMJ312.1(

nZR ⋅⋅−=⋅−= −

1s2 2s2 2p6 3s2 → 1s2 2s2 2p6 3s1 + e-

IE is the energy difference betweenthe total electron energy of these two configurations.

Slater’s Rules

1s2 2s2 2p6 3s2 → 1s2 2s2 2p6 3s1 + e-

DE = Σ (2E1f+8E2f+1E3f) - Σ (2E1i+8E2i+2E3i)

DE = 1E3f - 2E3i

= RH [-3.202/32] - 2 · RH [-2.852/32]

= 0.667 RH = 0.875 MJ/mol

The energy of inner e- areunaffected by outer e- shielding

Actual IE = 0.738 MJ/mol, 19% error

The basis of Slater’s ruless and p orbitals have better “penetration” to the nucleus than d (or f) orbitals for any given value of n

i.e. there is a greater probability of s and p electrons being near the nucleus

This means:

1. ns and np orbitals completely shield nd orbitals2. (n-1) s and p orbitals don’t completely shield n s and

p orbitals

Slater’s rules are only approximate and can give poor predictions. For example:

They ignore the differences in penetration between s and p orbitals. Real s and p orbitals do not have the same energy.

They assume that all electrons in lower shells shield outer electrons equally effectively.

Z* can be used to estimate ionization energy

Slater Calculation of Effective Nuclear Charge

0.00

1.00

2.00

3.00

4.00

5.00

6.00

7.00

8.00

0 5 10 15 20 25

Atomic Number

Zeff

Trends in Ionization EnergyThe amount of energy required to completely remove an electron from a gaseous atom.Removing one electron makes a 1+ ion.The energy required to remove the first electron is called the first ionization energy.

Ionization EnergyThe 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 or 2nd IE.

Symbol First Second ThirdHHeLiBeBCNOF Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

Symbol First Second ThirdHHeLiBeBCNOF Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

What determines IEThe greater the nuclear charge, the greater IE.Greater distance from nucleus decreases IEFilled and half-filled orbitalshave lower energy, so achieving them is easier, lower IE.Shielding effect

Group trends

As you go down a group, first IE decreases because...The electron is further away.More shielding.

Periodic trendsAll the atoms in the same period have the same energy level.Same shielding.But, increasing nuclear chargeSo IE generally increases from left to right.Exceptions at full and 1/2 full orbitals.

Firs

t Ion

izat

ion

ener

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Atomic number

He

He has a greater IE than H.same shielding greater nuclear charge

H

Firs

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ion

ener

gy

Atomic number

H

He

Li has lower IE than Hmore shielding further awayoutweighs greater nuclear charge

Li

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE than Lisame shielding greater nuclear charge

Li

Be

Firs

t Ion

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ion

ener

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Atomic number

H

HeB has lower IE than Besame shielding greater nuclear chargeBy removing an electron we make s orbital filled

Li

Be

B

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

Breaks the pattern, because removing an electron leaves 1/2 filled p orbital

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

NeNe has a lower IE than HeBoth are full,Ne has more shieldingGreater distance

Firs

t Ion

izat

ion

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Atomic number

H

He

Li

Be

B

C

N

O

F

NeNa has a lower IE than LiBoth are s1

Na has more shieldingGreater distance

Na

Firs

t Ion

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Atomic number

Driving Force

Full Energy Levels require lots of energy to remove their electrons.Noble Gases have full orbitals.Atoms behave in ways to achieve noble gas configuration.

2nd Ionization Energy

For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected.True for s2

Alkaline earth metals form 2+ ions.

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!!!

Electron Affinity

The tendency of an atom or anion to pick upexcess electron density is a measure of theelectron affinity of that atom.

F(g) + e– F– (g)

Electron AffinityAtoms which need a few electrons to attainan outer shell inert gas electronic configuration have high electron affinity.

Electron AffinityAtoms which tend to lose electrons to attainouter shell inert gas electronic configurationhave low electron affinity.

Trends in Electron AffinityEasiest to add to group 7A.

Gets them to full energy level.

Increase from left to right: atoms become smaller, with greater nuclear charge.

Decrease as we go down a group.

Trends in Atomic SizeFirst problem: Where do you start measuring from?The electron cloud doesn’t have a definite edge.They get around this by measuring more than 1 atom at a time.

Atomic Size

Atomic Radius = half the distance between two nuclei of a diatomic molecule.

Radius

Trends in Atomic Size

Influenced by three factors:1. Energy Level• Higher energy level is further away.2. Charge on nucleus• More charge pulls electrons in

closer.3. Shielding effect

Group trendsAs we go down a group...each atom has another energy level,so the atoms get bigger.

HLi

Na

K

Rb

Periodic TrendsAs you go across a period, the radius gets smaller.Electrons are in same energy level.More nuclear charge.Outermost electrons are closer.

Na Mg Al Si P S Cl Ar

Overall

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

Trends in Ionic SizeCations form by losing electrons.Cations are smaller than the atom they come from.Metals form cations.Cations of representative elements have noble gas configuration.

Ionic sizeAnions form by gainingelectrons.Anions are bigger that the atom they come from.Nonmetals form anions.Anions of representative elements have noble gas configuration.

Configuration of IonsIons 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.

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.

Periodic TrendsAcross the period, nuclear charge increases so they get smaller.Energy level changes between anions and cations.

Li1+

Be2+

B3+

C4+

N3-O2- F1-

Group trendsAdding energy levelIons get bigger as you go down.

Li1+

Na1+

K1+

Rb1+

Cs1+

Size of Isoelectronic ionsIso- means the sameIso electronic ions have the same # of electronsAl3+ Mg2+ Na1+ Ne F1- O2- and N3-

all have 10 electronsall have the configuration: 1s22s22p6

Size of Isoelectronic ionsPositive ions that have more protons would be smaller.

Al3+

Mg2+

Na1+ Ne F1- O2- N3-

ElectronegativityThe tendency for an atom to attract electrons to itself when it ischemically combined with another element.How fair is the sharing?Big electronegativity means it pulls the electron toward it.Atoms with large negative electron affinity have larger electronegativity.

Group Trend

The further down a group, the farther the electron is away, and the more electrons an atom has.More willing to share.Low electronegativity.

Periodic TrendMetals are at the left of the table.They let their electrons go easilyLow electronegativityAt the right end are the nonmetals.They want more electrons.Try to take them away from othersHigh electronegativity.

Ionization energy, Electronegativity, and Electron Affinity INCREASE

Atomic size increases

Ionic size increases

107:Bh 108: Hs 109: Mt 110: Ds107:Bh 108: Hs 109: Mt 110: Ds