PERIODICITYPERIODICITY Periodic Table Dmitri Mendeleev developed the modern periodic table. Argued that element properties are periodic functions of

PERIODICITYPERIODICITYPERIODICITYPERIODICITY

Periodic TablePeriodic Table• Dmitri Mendeleev developed the Dmitri Mendeleev developed the

modern periodic table. Argued modern periodic table. Argued that element properties are that element properties are periodic functions of their periodic functions of their atomic weights.atomic weights.

• We now know that element We now know that element properties are periodic functions properties are periodic functions of their of their ATOMIC NUMBERSATOMIC NUMBERS..

PeriodsPeriods in the Periodic Table in the Periodic Table

GroupsGroups in the Periodic Table in the Periodic Table

Regions of the Periodic Regions of the Periodic TableTable

Metals, Nonmetals, and Metalloids

– Three classes of elements are metals, nonmetals, and metalloids.

– Across a period, the properties of elements become less metallic and more nonmetallic.

6.1


»Metals, Metalloids, and Nonmetals in the Periodic Table

6.1



6.1



6.1



6.1


– Metals• Metals are good conductors of heat and electric

current.– 80% of elements are metals.

– Metals have a high luster, are ductile, and are malleable.

6.1


» Uses of Iron, Copper, and Aluminum

6.1



6.1



6.1


– Nonmetals• In general, nonmetals are poor conductors of heat

and electric current.– Most nonmetals are gases at room temperature.– A few nonmetals are solids, such as sulfur and

phosphorus.– One nonmetal, bromine, is a dark-red liquid.

6.1


– Metalloids• A metalloid generally has properties that are similar

to those of metals and nonmetals.

• The behavior of a metalloid can be controlled by changing conditions.

6.1

http://www.webelements.com/webelements/elements/text/Si/geol.html

Element Element AbundanceAbundance

FeC AlO Si

HydrogenHydrogenHydrogenHydrogen

Shuttle main engines use Shuttle main engines use HH22 and O and O22 The Hindenburg crash, The Hindenburg crash,

May 1939.May 1939.

Group 1A: Alkali Group 1A: Alkali MetalsMetals

Group 1A: Alkali Group 1A: Alkali MetalsMetals

Cutting sodium metalCutting sodium metal

Reaction of potassium + H2O

MagnesiumMagnesium

Magnesium Magnesium oxideoxide

Group 2A: Alkaline Earth MetalsGroup 2A: Alkaline Earth Metals

Calcium Carbonate—Calcium Carbonate—LimestoneLimestone

The Appian Way, ItalyThe Appian Way, Italy Champagne cave carved into Champagne cave carved into chalk in Francechalk in France

Group 3A: B, Al, Ga, In, Group 3A: B, Al, Ga, In, TlTl

Group 3A: B, Al, Ga, In, Group 3A: B, Al, Ga, In, TlTl

AluminumAluminum

Boron halidesBoron halides BF BF33 & BI & BI33

Gems & MineralsGems & Minerals

• Sapphire:Sapphire: Al Al22OO33

with Fewith Fe3+3+ or Ti or Ti3+3+ impurity gives impurity gives blue whereas Vblue whereas V3+3+ gives violet.gives violet.

• Ruby:Ruby: Al Al22OO33

with Crwith Cr3+3+ impurityimpurity

Group 4A: C, Si, Ge, Sn, Group 4A: C, Si, Ge, Sn, PbPb

Group 4A: C, Si, Ge, Sn, Group 4A: C, Si, Ge, Sn, PbPb

Quartz, SiOQuartz, SiO22

DiamondDiamond

Group 5A: N, P, As, Sb, BiGroup 5A: N, P, As, Sb, BiGroup 5A: N, P, As, Sb, BiGroup 5A: N, P, As, Sb, Bi

White and red White and red phosphorusphosphorus

Ammonia, NHAmmonia, NH33

PhosphoruPhosphoruss

• Phosphorus first isolated by Brandt from urine, 1669

Group 6A: O, S, Se, Te, PoGroup 6A: O, S, Se, Te, PoGroup 6A: O, S, Se, Te, PoGroup 6A: O, S, Se, Te, Po

Sulfuric acid dripping from Sulfuric acid dripping from snot-tite in cave in Mexicosnot-tite in cave in Mexico

Sulfur from Sulfur from a volcanoa volcano

Group 7A: Group 7A: F, Cl, Br, I, AtF, Cl, Br, I, AtGroup 7A: Group 7A:

F, Cl, Br, I, AtF, Cl, Br, I, At

Group 8A: Group 8A: He, Ne, Ar, Kr, Xe, RnHe, Ne, Ar, Kr, Xe, Rn

Group 8A: Group 8A: He, Ne, Ar, Kr, Xe, RnHe, Ne, Ar, Kr, Xe, Rn

• Lighter than air balloons

• “Neon” signs

XeOFXeOF44XeOFXeOF44

Transition ElementsTransition ElementsTransition ElementsTransition Elements

Lanthanides and actinidesLanthanides and actinides

Iron in air gives Iron in air gives iron(III) oxideiron(III) oxide

LithiumLithiumLithiumLithium

Group 1AGroup 1AAtomic number = 3Atomic number = 31s1s222s2s11 ---> 3 total electrons ---> 3 total electrons

1s

2s

3s3p

2p

BerylliumBerylliumBerylliumBeryllium

Group 2AGroup 2A

Atomic number = 4Atomic number = 4

1s1s222s2s22 ---> 4 total ---> 4 total electronselectrons

1s

2s

3s3p

2p

BoronBoronBoronBoron

Group 3AGroup 3AAtomic number = 5Atomic number = 51s1s2 2 2s2s2 2 2p2p11 ---> ---> 5 total electrons5 total electrons

1s

2s

3s3p

2p

CarbonCarbonCarbonCarbonGroup 4AGroup 4AAtomic number = 6Atomic number = 61s1s2 2 2s2s2 2 2p2p22 ---> ---> 6 total electrons6 total electrons

1s

2s

3s3p

2p

NitrogenNitrogenNitrogenNitrogen


1s

2s

3s3p

2p

OxygenOxygenOxygenOxygen


1s

2s

3s3p

2p

FluorineFluorineFluorineFluorine


1s

2s

3s3p

2p

NeonNeonNeonNeon


1s

2s

3s3p

2p

Colors of Transition Colors of Transition Metal CompoundsMetal Compounds

Iron Cobalt Nickel Copper Zinc

PERIODIC

PERIODIC TRENDS

TRENDS

PERIODIC

PERIODIC TRENDS

TRENDS

General Periodic TrendsGeneral Periodic Trends• Atomic and ionic sizeAtomic and ionic size• Ionization energyIonization energy• Electron affinityElectron affinity• ElectronegativityElectronegativity

Higher effective nuclear chargeElectrons held more tightly

Larger orbitals.Electrons held lesstightly.

Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)

• Z* is the nuclear charge experienced by the outermost electrons.Z* is the nuclear charge experienced by the outermost electrons. Is the result of the nuclear attraction being blocked by the core Is the result of the nuclear attraction being blocked by the core electrons. Nuclear attraction increases with an increase in protonselectrons. Nuclear attraction increases with an increase in protons

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

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

Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*• Shielding effect remains constant across a period. Shielding effect remains constant across a period.

As the nuclear attraction increases across the As the nuclear attraction increases across the shielding effect is less effective.shielding effect is less effective.

• Shielding effect increases down a group thus Shielding effect increases down a group thus effectively blocking any increase in nuclear effectively blocking any increase in nuclear attraction.attraction.

• Electrons with a higher quantum number have more Electrons with a higher quantum number have more kinetic energy and thus are less affected by the kinetic energy and thus are less affected by the nuclear charge.nuclear charge.

Each of these forces need to be accounted for in Each of these forces need to be accounted for in each trend.each trend.

EffectiveEffective Nuclear Charge, Nuclear Charge, Z*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

Periodic Trend in Periodic Trend in the Reactivity of the Reactivity of

Alkali Metals Alkali Metals with Waterwith Water

Periodic Trend in Periodic Trend in the Reactivity of the Reactivity of

Alkali Metals Alkali Metals with Waterwith WaterLithiumLithium

SodiumSodium PotassiumPotassium

Atomic Atomic SizeSize

Atomic Atomic SizeSize

• Size goes UPSize goes UP on going down a on going down a group. group.

• Because electrons are added Because electrons are added further from the nucleus, there is further from the nucleus, there is less attraction, due to an increase less attraction, due to an increase in sheilding effectiveness and in in sheilding effectiveness and in increase in kinetic energy.increase in kinetic energy.



Atomic SizeAtomic SizeAtomic SizeAtomic Size





General Outline for Trends

• Trend-define• Down a group• Nuclear attraction-define

once– Trend, effect

– Shielding effect-define once• Trend, effect

– Kinetic energy-define once• Trend, effect

• Across a period• Nuclear attraction

– Trend, effect

– Shielding effect• Trend, effect

– Kinetic energy• Trend, effect

Atomic Radius• Atomic radius is the distance from the nucleus to

the valance electrons. – Nuclear attraction (the attraction of the protons in the

nucleus on valance electrons) increases going down a group. This should pull the electrons in closer to the nucleus.

– Shielding effect (the blocking of nuclear attractions by core electrons) Shielding effect increases down a group offsetting the increase in nuclear attraction.

– Kinetic energy (the energy of valance electrons associated with principle energy levels) increases down a group allowing the valance electrons to orbit farther from the nucleus increasing atomic radius.

Atomic Radius

– Nuclear attraction increases across a period. This should pull the electrons in closer to the nucleus decreasing atomic radius.

– Shielding effect remains constant across a period not offsetting nuclear attraction.

– Kinetic energy remains constant across a period so effective nuclear attraction is greater and the atomic radius decreases.

Atomic RadiiAtomic RadiiAtomic RadiiAtomic Radii

Figure 8.9Figure 8.9

Atomic SizeAtomic SizeAtomic SizeAtomic SizeSize 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

Ion SizesIon SizesIon SizesIon Sizes

Li,152 pm3e and 3p

Li+, 60 pm2e and 3 p

+Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?


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

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


F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?

Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?


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

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

-

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

Figure 8.13Figure 8.13

Ionic Size• Ionic size is the distance from the nucleus to the

valence electrons after an atom has lost or gained electrons.

• Cations form when an atom loses one or more electrons.

• Cations are smaller than the atoms from which they form

• Ionic size - Cations– Effective nuclear charge increases dramatically when

electrons are removed.– Shielding effect decreases compared to the atom

because valence electrons are lost and some or all of the core electrons become valence electrons.

– Kinetic energy the new valence electrons have less kinetic energy to resist the pull of the nucleus.

Ionic Size• Anions form when an atom gains one or more

electrons• Ionic size - Anions

– After the addition of valence electron(s) the nuclear attraction is diluted.

– Shielding effect remains still resisting the nuclear attraction.

– Kinetic energy increases because of additional repulsion due to more electrons in the valence shell increasing the anions size.

– Anions are larger than the atoms from which they form

Ionization EnergyIonization EnergyIonization EnergyIonization Energy

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

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

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




MgMg2+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. shell of lower n.


General Periodic General Periodic TrendsTrends• Atomic and ionic sizeAtomic and ionic size

• Ionization energyIonization energy• Electron affinityElectron affinity

Higher Z*.Electrons heldmore tightly.

Larger orbitals.Electrons held lesstightly.

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

Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy• IE increases across a period IE increases across a period

because Z* increases.because Z* increases.• Metals lose electrons more Metals lose electrons more

easily than nonmetals.easily than nonmetals.• Metals are good reducing Metals are good reducing

agents.agents.• Nonmetals lose electrons with Nonmetals lose electrons with

difficulty.difficulty.

Trends in Ionization Trends in Ionization EnergyEnergy

Trends in Ionization Trends in Ionization EnergyEnergy

• IE decreases down a group IE decreases down a group

• Because size increases.Because size increases.

• Reducing ability generally Reducing ability generally increases down the periodic increases down the periodic table. table.

• See reactions of Li, Na, KSee reactions of Li, Na, K



MgMg2+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. shell of lower n. This is why ox. no. = Group no.This is why ox. no. = Group no.

Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12

Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12

Ionization Energy• Ionization energy is the energy

needed to remove an electron from an atom.

• Nuclear attraction

– Increases down a group holding the electrons tighter.

– Shielding effect increases down a group offsetting the increase in nuclear attraction.

– Kinetic energy increases down a group giving the electrons greater initial energy. This reduces the additional energy needed to remove an electron.

• Nuclear attraction increases across a period holding the electrons tighter.

• Shielding effect is constant across and does not offset the increase in nuclear attraction.

• Kinetic energy remains constant across. With the same initial energy valence electrons are increasingly harder to remove due to the greater effective nuclear charge.

Electron AffinityElectron Affinity

A few elements A few elements GAINGAIN electrons to electrons to form form anionsanions..

Electron affinity is the energy Electron affinity is the energy involved when an atom gains an involved when an atom gains an electron to form an anion.electron to form an anion.

A(g) + e- ---> AA(g) + e- ---> A--(g) E.A. = (g) E.A. = ∆E∆E

Electron Affinity of OxygenElectron Affinity of Oxygen

∆∆E is E is EXOEXOthermic thermic because O has because O has an affinity for an an affinity for an e-.e-.

[He] O atom

EA = - 141 kJ

+ electron

O [He] - ion

Electron Affinity of Electron Affinity of NitrogenNitrogen

∆∆E is E is zero zero for Nfor N- -

due to electron-due to electron-electron electron repulsions.repulsions.

EA = 0 kJ

[He] N atom

[He] N- ion

+ electron

• See Figure 8.12 and See Figure 8.12 and Appendix FAppendix F

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

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

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 AffinityTrends in Electron Affinity

Trends in Electron AffinityTrends in Electron Affinity

Electronegativity Values

See page 177 in text

Electronegativity• Electronegativity is the tendency

of an atom to remove an electron from another atom when forming a compound.

• Nuclear attraction

– Increases down a group attracting the electrons more.

– Shielding effect increases down a group offsetting the increase in nuclear attraction.

– Kinetic energy increases down a group giving the atom a larger radius and increasing the proximity of the nucleus to adjacent electrons decreasing electronegativity

• Nuclear attraction increases across a period attracting electrons more.

• Shielding effect and kinetic energy are constant across. This increases effective nuclear charge allowing the atom to remove electron from other atoms with lesser electronegativity.

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PERIODICITYPERIODICITY Periodic Table Dmitri Mendeleev developed the modern periodic table. Argued that element properties are periodic functions of