1

PERIODIPERIODIC C

TRENDSTRENDS

PERIODIPERIODIC C

TRENDSTRENDS

2

Electron Electron Filling Filling OrderOrder

Figure 8.5Figure 8.5

3

Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table

Figure 8.7

4

PhosphorusPhosphorusPhosphorusPhosphorus

All Group 5A All Group 5A elements have elements have [core ] ns[core ] ns2 2 npnp3 3

configurations configurations where n is the where n is the period number.period number.

Group 5AGroup 5A

Atomic number = 15Atomic number = 15

1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p33

[Ne] 3s[Ne] 3s2 2 3p3p33

1s

2s

3s3p

2p

5

LithiumLithiumLithiumLithium

Group 1A

Atomic number = 3

1s22s1 ---> 3 total electrons

1s

2s

3s3p

2p

6

Electron PropertiesElectron Properties

Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

7

NeonNeonNeonNeon

Group 8A

Atomic number = 10

1s2 2s2 2p6 --->

Diamagnetic

1s

2s

3s3p

2p

8

BerylliumBerylliumBerylliumBeryllium

Group 2A

Atomic number = 4

1s22s2

Diamagnetic

1s

2s

3s3p

2p

9

BoronBoronBoronBoron

Group 3A

Atomic number = 5

1s2 2s2 2p1

Paramagnetic

1s

2s

3s3p

2p

10

CarbonCarbonCarbonCarbon

Group 4A

Atomic number = 6

1s2 2s2 2p2

Paramagnetic1s

2s

3s3p

2p

11

FluorineFluorineFluorineFluorine

Group 7A

Atomic number = 9

1s2 2s2 2p5 --->

Paramagnetic

1s

2s

3s3p

2p

12

Ion ConfigurationsIon ConfigurationsIon ConfigurationsIon Configurations

How do we know the configurations of ions?

Determine the magnetic properties of ions.

Ions with UNPAIRED ELECTRONS are

PARAMAGNETIC.

Without unpaired electrons DIAMAGNETIC.

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transition metal ionstransition metal ionstransition metal ionstransition metal ions

Fe [Ar] 4s2 3d6 loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6

• loses 3 electrons ---> Fe3+ [Ar] 4s0 3d5

4s 3d 3d4s

Fe Fe2+

4s 3d 3d4s

Fe Fe2+

14

Transition MetalsTransition Metals

How do they fill? How can we determine?

CopperCopperIronIron

ChromiumChromium

15

Ion Configurations MnIon Configurations MnIon Configurations MnIon Configurations Mn

Mn [Ar] 4s2 3d5 ---> Mn5+ [Ar] 4s03d2

loses 5 electrons ---> Mn5+ [Ar] 4s2 3d0

4s 3d 3d4s

Fe Fe2+D P

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PERIODIPERIODIC C

TRENDSTRENDS

PERIODIPERIODIC C

TRENDSTRENDS

17

PERIODICITYPERIODICITYPERIODICITYPERIODICITYPeriod Law-

-physical and chemical properties of elements are a

periodic function of their atomic numbers

18General Periodic General Periodic TrendsTrends

• Atomic and ionic size

• Ionization energy

• Electron affinity, electronegativity

19Effective Nuclear Effective Nuclear ChargeCharge

Z*Z*

20

Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*

Effective Nuclear Charge, Effective Nuclear Charge, Z*Z*

• Z* is the nuclear charge experienced by Z* is the nuclear charge experienced by the outermost electrons.the outermost electrons. See p. 295 and Screen 8.6.See p. 295 and Screen 8.6.

• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)

• Z* increases across a period owing to Z* increases across a period owing to incomplete shielding by inner electrons.incomplete shielding by inner electrons.

• Estimate Z* by --> [ Estimate Z* by --> [ Z - (no. inner electrons) Z - (no. inner electrons) ]]

• Charge felt by 2s e- in Li Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Z* = 3 - 2 = 1

• Be Be Z* = 4 - 2 = 2Z* = 4 - 2 = 2

• B B Z* = 5 - 2 = 3Z* = 5 - 2 = 3 and so on!and so on!

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EffectiveEffective Nuclear Charge, Z* Nuclear Charge, Z*

• Atom Z* Experienced by Electrons in Valence Orbitals

• Li +1.28

• Be -------

• B +2.58

• C +3.22

• N +3.85

• O +4.49

• F +5.13

Increase in Increase in Z* across a Z* across a periodperiod

22General Periodic General Periodic TrendsTrends

Higher effective nuclear chargeElectrons held more tightly

Larger orbitals.

Electrons held less

tightly.

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Periodic Trend in the Reactivity of

Metals

Periodic Trend in the Reactivity of

Metals

LithiumLithium

SodiumSodium

PotassiumPotassium

MOST

2. Reactivity for Metals

As you go down a group for metals the number of energy levels increase.

Because of this, reactivity increases because the atom is more willing to give away its electron (react).

3.Nonmetalic Trends: Gain electrons Nonmetals on right side, form anions

Going right elements are more nonmetallic (better gainers of electrons)

Going UP elements become more nonmetallic (want to gain)

8. Reactivity nonmetals: Gain e The reason Across = fill the energy level

Going UP a group, nonmetals have same valence but fewer total electrons

Flourine is the most reactive nonmetal.

27Atomic Atomic RadiiRadiiAtomic Atomic RadiiRadii

Figure 8.9Figure 8.9

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Atomic Atomic SizeSize

Atomic Atomic SizeSize

• Size increasesSize increases, down a group. , down a group.

• Because electrons are added into Because electrons are added into additional energy levels, there is less additional energy levels, there is less attraction.attraction.

• Size decreases Size decreases across a period.across a period.

• Because, increased effective nuclear Because, increased effective nuclear charge.charge.

• Size increasesSize increases, down a group. , down a group.

• Because electrons are added into Because electrons are added into additional energy levels, there is less additional energy levels, there is less attraction.attraction.

• Size decreases Size decreases across a period.across a period.

• Because, increased effective nuclear Because, increased effective nuclear charge.charge.

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Atomic SizeAtomic SizeAtomic SizeAtomic Size

Size Size decreasesdecreases across a period owing across a period owing to increase in Z*. Each added electron to increase in Z*. Each added electron feels a greater and greater + charge.feels a greater and greater + charge.

LargeLarge SmallSmall

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Trends in Atomic SizeTrends in Atomic SizeSee Figures 8.9 & 8.10See Figures 8.9 & 8.10

0

50

100

150

200

250

0 5 10 15 20 25 30 35 40

Li

Na

K

Kr

He

NeAr

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

0

50

100

150

200

250

0 5 10 15 20 25 30 35 40

Li

Na

K

Kr

He

NeAr

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

31

Ion SizesIon SizesIon SizesIon Sizes

• CATIONSCATIONS are are SMALLERSMALLER than the than the atoms from which they come.atoms from which they come.

• The electron/proton attraction has The electron/proton attraction has gone UP and so size gone UP and so size DECREASESDECREASES..

Li,152 pm3e and 3p

Li +, 78 pm2e and 3 p

+Forming Forming a cation.a cation.Forming Forming a cation.a cation.

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Ion SizesIon SizesIon SizesIon Sizes

• ANIONSANIONS are are LARGERLARGER than the atoms from than the atoms from which they come.which they come.

• The electron/proton attraction has gone DOWN The electron/proton attraction has gone DOWN and so size and so size INCREASESINCREASES..

• Trends in ion sizes are the same as atom sizes. Trends in ion sizes are the same as atom sizes.

Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm

9e and 9pF-, 133 pm10 e and 9 p

-

33

Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes

Figure 8.13Figure 8.13

34

Ionization EnergyIonization Energy

IE = energy required to remove an electron from an atom in the gas phase.

Mg (g) + 738 kJ ---> Mg+ (g) + e-

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Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-

MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-

MgMg++ has 12 protons and only 11 has 12 protons and only 11 electrons. Therefore, IE for Mgelectrons. Therefore, IE for Mg++ > Mg. > Mg.

IE = energy required to remove an electron from an atom in the gas phase.

Ionization EnergyIonization Energy

36

11stst IE: Mg (g) + 735 kJ ---> Mg IE: Mg (g) + 735 kJ ---> Mg++ (g) + e- (g) + e-

22ndnd IE: Mg IE: Mg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-

33rdrd IE: Mg IE: Mg2+2+ (g) + 7733 kJ ---> Mg (g) + 7733 kJ ---> Mg3+3+ (g) + e- (g) + e-

Energy cost is very high to dip into a Energy cost is very high to dip into a shell of lower n (core electrons). shell of lower n (core electrons). This is why ox. no. = Group no.This is why ox. no. = Group no.

Ionization EnergyIonization Energy

37

Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy

1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 350

500

1000

1500

2000

2500

1st Ionization energy (kJ/mol)

Atomic NumberH Li Na K

HeNe

ArKr

38

Trends in Ionization Trends in Ionization EnergyEnergy

Trends in Ionization Trends in Ionization EnergyEnergy

• IE decreases down a group

• Because size increases.

• IE increases across a period

• Because effective nuclear

charge increases

• IE increases across a period

• Because effective nuclear

charge increases

39

Electron AffinityElectron Affinity

A few elements GAIN electrons to form anions.

Electron affinity is the energy involved when an atom

gains an electron to form an anion.

X(g) + e- ---> X-(g) E.A. = ∆E

40

Trends in Electron AffinityTrends in Electron Affinity

41

• Affinity for electron increases across a period (EA becomes more negative).

• Affinity decreases down a group (EA becomes less negative).

Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Trends in Electron Trends in Electron AffinityAffinity

42

Electron Affinity of OxygenElectron Affinity of Oxygen

∆E is EXOthermic

because O has an affinity for

an e-.

[He] O atom

EA = - 141 kJ

+ electron

O [He] - ion

43

Electron Affinity of Electron Affinity of NitrogenNitrogen

∆∆E is E is zero zero for Nfor N- -

due to due to electron-electron-electron electron

repulsions.repulsions.EA = 0 kJ

[He] N atom

[He] N- ion

+ electron

44

Electronegativity

• So how is this different from electron affinity?

• Electron Affinity – is rating of how well an atom wants to gain an electron

• Electronegativity – is rating of how well an atom keeps the electron once it is bonded to another atom

45

Electronegativity

46Electron Configurations and the Periodic Trends

47

“Your Best Friend”

• Periodic table