1

Periodic Variation in Physical Properties

Sizes of atoms and ionsIonization energyElectron affinity

Metallic properties

2

Properties and Electronic Structure

Properties depend on

• valence electrons

• Effective nuclear charge (net charge on an electron)Protons-electrons attraction (increases)Electron-electron repulsion (decreases)Penetration of the orbitals (greater penetration –

increases the Zeff)

• The shielding effect Core electrons

• Size of the atom

3

Effective Nuclear Charge

In a many-electron atom,

the Zeff depends on two

factors:

1. Attraction between electrons and nucleus

2. Repulsion between electrons in orbitals

Effective Nuclear charge, Zeff: the actual nuclear charge a valence electron experiences.

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Effective Nuclear Charge

In SWE, electron is treated individually in a field of net nuclear charge that determines the Zeff..

The effective nuclear charge, Zeff, is found:

Zeff = Z − σ Z =atomic number

σ =screening constant~ # number of inner electrons, but not equal.

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Effective Nuclear Charge (Zeff)

Proton-electron attraction increases the Zeff.

Size of the atom: Coulomb’s Law

Penetration of orbitals: (10% probability being next to the nucleus)

It depends on:

6

Effective Nuclear Charge (Zeff)

Electron-electron repulsion reduces the Zeff

Smaller Zeff on valence electron, easier to remove the electron from the atom

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Effective Nuclear Charge on Valence Electrons

• Expected Nuclear Charge by the valence electron in – Li atom; Be atom; B atom …..

But: Li. Zeff on the 3rd electron 1.3. Why?

Zeff = Z – (shielding by other electrons)

Zeff = Z – σ

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Effective Nuclear Charge on 2s Electron in Li

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Shielding and (Zeff) in a Group

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•The electron on the outside energy level has to look through all the other energy levels to see the nucleus: it is shielded from the nucleus by all the inner electrons

•Less attraction between the valence electron and the nucleus

•Lower effective nuclear charge

•If shielding were completely effective, Zeff = 1

•Why isn’t it?

10

Zeff in a Group

• Electrons enter new shell (energy level)• Size of atom increases• Zeff decreases: decreased force of attraction (Coulomb’s

Law – inversely proportional to r2)

• In a Group: Zeff decreases; shielding Increases.

11

Shielding in a Group

• The more positive the nucleus (higher n, higher atomic number), the closer the inner orbitals to the nucleus

provides for better shielding of the outer electrons (easier to remove electron in Na than in Li or H)

by the core electrons.

• In lithium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus.

• The sodium 1s is even smaller.

12

Shielding: Electron Repulsion and Orbital Energy

Extent of shielding depends on:

1. Core e- provide more shielding than valence electrons

2. Electrons in the same shell ( same n), extend of shielding depends on l (penetration effects);

s > p > d > f

Correlates with degree of penetration.

3. Electrons in the same sub-shell (same value of n and l) do not effectively shield one another. Electrons are in the degenerate orbitals

Shielding in Period

• The electron on the outside energy level has to look through all the other energy levels to see the nucleus

• A second electron in the same energy level has almost the same shielding.

• The third one has almost the same shielding

• Zeff (effective nuclear charge) increases as atomic number increases in the Period.

• Why?

29

14

Shielding Effects: Examples

First Ionization energies in kJ/mol

• He atom (1s2 + 2 protons): 2372• He+ ion: (1s1 + 2 protons): 5250 ( it is not 2X 2372, but

larger)) • In He, the second electron repels the first, making it

easier to remove (one electron shields the other from the full effect of the nucleus); not in He+1 (there is only one electron)

• In Li (1s22s1): 520; Li+2 (1s1): 2954• In Li: we have two 1s electrons. They shield very

effectively the electron in 2s and smaller radius, stronger attraction

15

Sizes of Atoms

The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

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Atomic Radii: Covalent Compounds

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Atomic Size in a Group (shielding dominates)

– Each new member has one more level of inner electrons

– Inner electrons shield the outer electrons very effectively

– Zeff increases very slightly (more protons)

– Atoms get larges, as n-increases

– Atoms radii increase in a group

18

Atomic Size in a Period: Zeff Dominates

– Electrons added to the same shell– shielding by inner electrons changes very slightly, if at all. – Outer electrons shield each other poorly

– Zeff rises significantly, electrons pulled closer to nucleus

– Atomic radius decreases in a period.

Na Mg Al C Si P Cl Ar

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Atomic Radii of the Main Group Elements

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Overall

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

VZn

Ga

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Radii in Transition Elements

• Size shrinks for the first two to three members because of increased nuclear charge

• After: the size remains relatively constant as repulsion of d-electrons (increased radius) counteracts the increase in Zeff.

• d-electrons shield very well, but p-orbital penetrates much more than d-orbital: thus Ga (135 pm) is much smaller than Ca (197 pm)

• Another anomaly: 13Al (143 pm) versus 31Ga (135 pm). Filling the d-orbitals causes major contraction.

22

Atomic Radius: Examples

• Using only the periodic table, rank each set of main-group elements in order of decreasing atomic size:

• A) Ca, Mg, Sr• B) K, Ga, Ca• C) Br, Rb, Kr• D) Sr, Ca, Rb

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Order the following according to increasing atomic radius.

1. Ge < Si < Se < Cl 2. Se < Si < Ge < Cl

3. Si < Cl < Ge < Se

4. Cl < Si < Se < Ge

5. Si < Ge < Se < Cl

Ge Si Se Cl

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Ionic Radius

25

Ionic Radius

Ionic radius:

the radius of a cation or an anion.

Determines the physical and chemical properties of an ionic compound such as– Crystal structure– Melting point– solubility

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Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

8.3

278.3

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Sizes of Ions

• Ionic size depends upon:– Nuclear charge.– Number of

electrons.– Orbitals in which

electrons reside.

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Sizes of Ions

• In an isoelectronic series, ions have the same number of electrons.

• Ionic size decreases with an increasing nuclear charge.• All have the configuration 1s12s22p6 (10 electrons)

Radii of Atoms and Corresponding Ions

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Ionic Radius: Examples

• In each of the following pairs, indicate which one of the two species is larger:

• A) N3- or F-

• B) Mg+2 or Ca+2

• C) Fe+2 or Fe+3

• Explain your choice.

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Order the following according to increasing atomic/ionic radius.

1. C < Li+ < O2- < N3-

2. N3- < O2- < C < Li+

3. Li+ < C < N3- < O2-

4. Li+ < C < N3- < O2- 5. Li+ < C < O2- < N3-

N3- Li+ C O2-

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Ionization Energy

Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.– First ionization energy is that energy required

to remove first electron.– Second ionization energy is that energy

required to remove second electron, etc.

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Ionization Energy

• It requires more energy to remove each successive electron.

• When all valence electrons have been removed, the ionization energy takes a quantum leap.

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Trends in First Ionization Energies

• As one goes down a column, less energy is required to remove the first electron.– For atoms in the same

group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

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Filled n=1 shell

Filled n=2 shell

Filled n=3 shellFilled n=4 shell

Filled n=5 shell

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Trends in First Ionization Energies

• Generally, as one goes across a row, it gets harder to remove an electron.– As you go from left to

right, Zeff increases, shielding is almost the same.

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Trends in First Ionization Energies

However, there are two apparent discontinuities in this trend.

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Trends in First Ionization Energies

• The first occurs between Groups IIA and IIIA.

• Electron removed from p-orbital rather than s-orbital– Electron farther from

nucleus– Small amount of

repulsion by s electrons.

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Trends in First Ionization Energies

• The second occurs between Groups VA and VIA.– Electron removed

comes from doubly occupied orbital.

– Repulsion from other electron in orbital helps in its removal.

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Trends in Successive IE

• The effective nuclear charge increases as you remove electrons. Thus IE3>IE2>IE1

• Big jump after all outer electrons removed.• It takes much more energy to remove a

core electron than a valence electron because there is less shielding, smaller size (energy shell removed)

• greater effective nuclear charge.

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Na Mg Al Si P S Cl Ar

IE1 496 738 578 787 1012 1000 1251 1520

IE2 4562 1451 1817 1577 1903 2251 2297 2665

IE3 6912 7733 2745 3231 2912 3361 3822 3931

IE4 9543 10540 11575 4356 4956 4564 5158 5770

IE5 13353 13630 14830 16091 6273 7013 6540 7238

IE6 16610 17995 18376 19784 22233 8495 9458 8781

IE7 20114 21703 23293 23783 25397 27106 11020 11995

Successive Ionization Energies (kJ/mol)  

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Explain the trend in IE

• For Mg • IE1 = 735 kJ/mole• IE2 = 1445 kJ/mole• IE3 = 7730 kJ/mole

• For Al• IE1 = 580 kJ/mole• IE2 = 1815 kJ/mole• IE3 = 2740 kJ/mole• IE4 = 11,600 kJ/mole

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Which will have the highest ionization energy?

• C• N• O• Al• Si

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Which will be the largest?

1. I1 of Na

2. I2 of Na

3. I1 of Mg

4. I2 of Mg

5. I3 of Mg

I = ionization energy

#5

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Electron Affinity

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Electron Affinity, kJ/mol

• Electron affinity: the energy change that occurs when an electron is accepted by an atom in gaseous state.

X(g) + e- → X-(g)

• A large negative value indicates a strong attraction between the atom and the added electron

Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol

• A positive value indicates the addition of electron is unfavorable

Ne(g) + e- → Ne- ΔE = 29 kJ/mol

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Trends in Electron Affinity

In general, electron affinity becomes more exothermic as you go from left to right across a row.

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Trends in Electron Affinity

There are again, however, two discontinuities in this trend.

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Trends in Electron Affinity

• The first occurs between Groups IA and IIA.– Added electron must

go in p-orbital, not s-orbital.

– Electron is farther from nucleus and feels repulsion from s-electrons.

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Trends in Electron Affinity

• The second occurs between Groups IVA and VA.– Group VA has no

empty orbitals.– Extra electron must

go into occupied orbital, creating repulsion.

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Summary

• Reactive nonmetals (Groups 6A and 7A): have high ionization energies and high negative electron affinities. Gain electrons easily, lose electrons with difficulties.

• Reactive metals (Groups 1A and 2A) have low ionization energies and slightly negative (exothermic) electron affinities. Lose electrons lightly and gain electrons with difficulties.

• Noble Gases: very high IE and slightly positive electron affinities. Do not lose or gain electrons easily.

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Metallic Character

• Metals– malleable & ductile– shiny, lustrous– conduct heat and

electricity– most oxides basic

and ionic– form cations in

solution– lose electrons in

reactions - oxidized

• MetalloidsAlso known as

semi-metalsShow some

metal and some nonmetal properties

• Nonmetalsbrittle in solid statedullelectrical and

thermal insulatorsmost oxides are

acidic and molecular form anions and

polyatomic anionsgain electrons in

reactions - reduced

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General Trends in Chemical Reactivity: DIAGONAL RULE

• First member in each group differs from the rest of the group

• It resembles element to its right and next period (Li-Mg; Be-Al, B-Si). Has to do with size. Called diagonal relationship.

55

Properties of Metal, Nonmetals,and Metalloids

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Metals versus Nonmetals

Differences between metals and nonmetals tend to revolve around these properties.

57

Metallic Character

Metals• Lose electrons to become positive ions• Low ionization energy• Low electron affinity• Good reducers

Nonmetals• High IE• Low EA• Gain (to become negative ions)• or share electrons to form covalent compounds• Oxidizers

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Metals versus Nonmetals

• Metals tend to form cations.• Nonmetals tend to form anions.

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Metals

Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

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Metals

• Compounds formed between metals and nonmetals tend to be ionic.

• Metal oxides tend to be basic.

• Form bases when reacted with water.

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Nonmetals

• Dull, brittle substances that are poor conductors of heat and electricity.

• Tend to gain electrons in reactions with metals to acquire noble gas configuration.

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Nonmetals

• Substances containing only nonmetals are molecular compounds.

• Most nonmetal oxides are acidic.

• Form acids with water.

63

Metalloids

• Have some characteristics of metals, some of nonmetals.

• For instance, silicon looks shiny, but is brittle and fairly poor conductor.

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Group Trends

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Alkali Metals

• Soft, metallic solids.• Name comes from

Arabic word for ashes.

66

Alkali Metals

• Found only as compounds in nature.• Have low densities and melting points.• Also have low ionization energies.

67

Alkali Metals

Their reactions with water are famously exothermic.

68

Alkali Metals

• Alkali metals (except Li) react with oxygen to form peroxides.

• K, Rb, and Cs also form superoxides:

K + O2 KO2

• Produce bright colors when placed in flame.

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Alkaline Earth Metals

• Have higher densities and melting points than alkali metals.

• Have low ionization energies, but not as low as alkali metals.

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Alkaline Earth Metals

• Be does not react with water, Mg reacts only with steam, but others react readily with water.

• All react with acids• Reactivity tends to

increase as go down group.

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Group 3A Elements (ns2np1, n≥2)

• B is a metalloid, does not react with oxygen or water; does not form binary ionic compounds

• Rest of elements mostly behave as metals

2Al(s) + 6H+(aq) → 2 Al+3(aq) + 3H2(g)

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Group 4A/5A Elements

• Group VA:– Carbon: nonmetal– Silicon and germanium: metalloids– Rest (Pb and Sn) metallic: react with acids,

but not with water.– Variable oxidation states for Pb and Sn.

• Group 5A:– N, P, nonmetals; As, and Sb metalloids– Rest are metals. P exists as P4. N2 forms

many oxides, P only 2.

73

Group 6A

• Oxygen, sulfur, and selenium are nonmetals.• Oxygen – strong oxidizer.• Tellurium is a metalloid.• The radioactive polonium is a metal.

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Oxygen

• Two allotropes:– O2

– O3, ozone

• Three anions:– O2−, oxide– O2

2−, peroxide– O2

1−, superoxide

• Tends to take electrons from other elements (oxidation)

75

Sulfur

• Weaker oxidizing agent than oxygen.

• Most stable allotrope is S8, a ringed molecule.

76

Group VIIA: Halogens

• Prototypical nonmetals• Name comes from the Greek halos and gennao:

“salt formers”

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Group VIIA: Halogens

• Large, negative electron affinities– Therefore, tend to oxidize

other elements easily

• React directly with metals to form metal halides

• Chlorine added to water supplies to serve as disinfectant

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Group VIIIA: Noble Gases

• Astronomical ionization energies• Positive electron affinities

– Therefore, relatively unreactive

• Monatomic gases

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Group VIIIA: Noble Gases

• Xe forms three compounds:– XeF2

– XeF4 (at right)

– XeF6

• Kr forms only one stable compound:– KrF2

• The unstable HArF was synthesized in 2000.

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Hydrogen: in a class by itself

• Acts as a metal: hydrated ion in solution• Acts as a nonmetal: forms hydrides

NaH(s) + H2O(l) → 2NaOH(aq) + H2(g)

NaH(s) + H2O(l) → 2Na+(aq) + OH-(aq) + H2(g)

Most important reaction:

2H2(g) + O2(g) → 2H2O(l)

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Which will produce a basic solution in water?

1. CO2

2. P2O5

3. BaO

4. XeO3

5. SO2

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Periodic Trends Interactive

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