_the chemistry of element

7/25/2019 _The Chemistry of Element

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CHM 474: INORGANIC CHEMISTRY I

The Chemistry of the Elements

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Chapter outline

-To know the periodic classification of the elements-Able to construct its electron configuration

-Identify the periodic variation in physical properties:

Effective nuclear charge (Zeff)

Atomic & ionic radii

Ionization Energy

Electron Affinity

- Identify the periodic variation in chemical properties

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Introduction.

Dmitri Mendeleev-Work on periodic

classification of elements according t o

their properties..

-Most significant achievement in 19th

century..

ns

1

ns

2

ns

2np

1

ns

2np

2

ns

2np

3

ns

2np

4

ns

2np

5

ns

2np

6

d1

d5

d10

4f

5f

Ground State Electron Configurations of the Elements

PERIODIC CLASSIFICATION OF THE ELEMENTS

n= principal quantumnumber oftheoutermost subshell

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Classification of the Elements

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Continued.Group 1A 7A (known as representative elements/main

group elements) Have incompletely filled sorpsubshells

Group 8A (except He) Show completely filled p subshells.Example:-He 1s2 and ns2np6 for other noble gases (Ne: 1s2

2s2 2p6)

Group 3B 8BKnown as d-block transition elements. Haveincompletely filledd subshells.They will produce cations with these incompletely filled d

subshells.

Group 4F & 5F Known as f block transition elements.

Have incompletely filled fsubshellsLanthanides Actinides

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Ground state electronicconfigurations (it is an electronicarrangement described for each

atom)

The aufbau PrincipleAufbau means building up

Used together with Hunds rules andPauli exclusion principle

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In the ground state of the atom,electrons will occupy the lowestenergy orbitals first, and onlyfill the higher energy orbitalswhen no lower energy orbitalsare left.

Hunds first rule:- electronsoccupy all the orbitals of a givensubshell singly before pairingbegins. These unpaired electronshave parallel spins.

Pauli exclusion principle:- no twoelectrons in the same atom mayhave the same set of n, l, ml, msquantum numbers

Fig. Order for filling

energy sublevels

with electrons.


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Figure 8.4

Condensed ground-state electron configurations inthe first three periods

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Electron Configurations of Cations and Anions

Na [Ne]3s1

Na+

[Ne]Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s 23p1 Al3+ [Ne]

Atoms lose electrons sothat cation has a stablenoble-gas outer electronconfiguration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or

[Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electronsso that anion has astable noble-gas outerelectron configuration.

Of Representative Elements

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+1

+2

+3

-1-

2-3

Cations and Anions of Representative Elements.

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Na+: [Ne]

Al3+: [Ne]

F-: 1s22s22p6 or [Ne]

O2-

: 1s2

2s2

2p6

or [Ne]N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

What neutral atom is isoelectronic with H- ?

Answer: H-: 1s2 same electron configuration as He

They have the samenumber of electronsand ground state

electronconfiguration

Quiz 2:

Isoelectronic

Examples:-

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Electron Configurations of Cations of Transition Metals

When a cation is formed from an atom of a transition

metal,electrons are always removed first from the nsorbital and then from the (n 1)d orbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6 or [Ar]3d6

Fe3+: [Ar]4s03d5 or [Ar]3d5

Examples:-

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5 or [Ar]3d5

Mn2+: [Ar]4s23d3

3d orbital is more stablethan the 4s orbital intransition metal ions..

1

.. 2

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ALL Periodic Table Trends

In f luenced by three fac tors :1.Energy Level

Higher energy levels are further away fromthe nucleus.

2.Charge on nucleus(# protons)

More charge pulls electrons in closer. (+ and attract eachother)

3.Shielding effect

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Shielding Effect

The electron on the outermost

energy level has to look through

all the other energy levels to seethe nucleus.

Second electron has sameshielding, if it is in the same

period

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Atomic Size

Measure the Atomic Radius- this is half the distance between the twonuclei of adiatomic molecule.

}

Radius

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What do they influence?

Energy levelsand Shieldinghave an effect on the

GROUP

Nuclear chargehas an effect on aPERIOD

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Periodic Variation in Physical Properties

A. Effec tive nuclear charge (Zeff) It is the positi ve charge felt by an electron

Given by Zeff= Z s

Z= actual nuclear charge,

s

= shielding/screening constant

(0 < s < Z)

ZeffZ number of inner or core electrons

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius

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increasing Zeff

increasingZ

eff Li Be B C N O F Ne

3 4 5 6 7 8 9 10Core electron = 1s2

1. Core electron > closer to nuclues than valence electron, thus core e shieldvalence e> than valence e shield each other.

2. Moving across the period, core e remains constant , but Z increases.3. The added e will be valence e, and due to valence e does not shield each other,

thus, moving across the period, > Zeff will be felt by valence e.4. Moving the group, Zeff . As n increases, large shells increases, thus valence e

are added to these large shells. Thus, electrostatic attraction between nuclues& valence e decreases.

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Example:

Which elements outer shell or valenceelectrons is predicted to have the largestEffective nuclear charge? Cl, O, N or Ca?

Cl: Zeff 17 - 10 = 7

O: Zeff 8 - 2 = 6

N: Zeff 7 - 2 = 5

Ca: Zeff 20 - 18 = 2


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Atomic Size - Group trends

Increasesdown agiven group.

the number of electrons and filled

electron shells increases

electrons are found further fromthe nucleus

Therefore, the atomic radiiincrease.

H

Li

Na

K

Rb

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#1. Atomic Size - Period Trends

Decreasesacross aperiod fromleft to right

Electrons are in the same energy level. and unshielded towardsattraction by protons.

protons are being added to the nucleus thus creates a"highereffective nuclear charge."

stronger force of attraction pulling the electrons closer to the

nucleus

Na Mg Al Si P S Cl Ar 22

Atomic Radii

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Ion ic Rad i i

1) Anions (negative ions) are larger than their respective atoms.

WHY? Electron-electron repulsion forces themto spread further apart.

the protons cannot pull the extraelectrons as tightly toward the nucleus.

2) Cations(positive ions) are smaller than their respectiveatoms.

WHY?There is less electron-electron repulsion.

Protons outnumber electrons; the protons can pull the fewer electrons towardthe nucleus more tightly.

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Ion Group trends

Eac h st ep dow n a g roup i s

ad d ing an en ergy lev e l .

I ons t he re f o re ge t b igger asy ou go d ow n, bec ause o f

t he add i t i ona l energy lev e l

Li1+

Na1+

K1+

Rb1+

Cs1+

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Ion Period Trends

Across the period fromleft to right, the nuclear charge increases -

so they get smaller.

Notice theenergy level changesbetween anions and cations.

Li1+

Be2+

B3+

C4+

N3-

O2- F1-

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QUIZ 4:

Arrange the following atoms in order ofincreasing atomic radius P, Si, N

STRATEGY: From left to right across periodDecreases

Moving up to down the group - Increases

ANSWER: N


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Positive ions that have more protons would besmaller(moreprotons would pull the same no. of electrons in closer)

Size of Isoelectronic ions?

Al3+ Mg2+

Na1+Ne F

1- O2-N3-

13 12 1110 9 8 7

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QUIZ 5:

Which one of two species is larger?????(a)N 3- or F-

(b)Mg2+ or Ca2+

STRATEGY: Think whether they are isoelectronic ions or are they

from the same group or period?????

ANSWER:(a) N3-

(both are isoelectronic and have 10 electrons.However, N3- (7 protons), F-(9 protons)thus lessattraction exerted by nuclues on the electrons in

larger N3-

(b) Ca2+

Both are in Group 2A. Ca in larger shell (n =4)Mg in n = 3.

Ionization energy is the minimum energy (kJ/mol)required to remove an electron from a gaseous atom in its

ground state.

- Depends on how tightly the is held in the atom

I1 + X (g) X+

(g) + e-

I2 + X (g) X2+

(g) + e-

I3 + X (g) X3+

(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ioni zation energy

I1 < I2 < I3

C. Ionization Energy

Higher t he i on i za t i on ene rgy ,the > d i f f i cu l t t o rem ove the

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1)Down a group, firstionization energy decreases.

Electrons are further fromthe nucleus

more shielding

easier to remove the outermost electron

2)Across a period, first ionization energy increases.

the atomic radius decreases

The outer electrons are closer to the nucleus and more

strongly attracted to the center

Similar shielding effect

more difficult to remove the outermost electron.

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FirstIonization

En

ergy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne

Na

First IonizationEnergy (IE) Trends

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Exceptions to First Ionization EnergyTrends

1) Xs2> Xp1

Example :4Be >5B

The energy of an electron in an Xp orbital is greater than Xs orbital.

less energy to remove the first electron in ap orbital than it is to remove onefromafilled s orbital.

2) Xp3> Xp4

example : 7N >8O

After the separate degenerate orbitals have been filled with single electrons,the fourth electron must be paired. The electron-electron repulsion makes iteasier to remove the outermost, paired electron. (Hund's Rule)

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Second and Higher Ionization Energies

Definition: Second Ionization Energy is the energy required to

remove asecond outermost electron fromaground state atom.

Subsequent ionization energiesincreasegreatly once an ion hasreached the state like that of anoble gas.

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

(True for s2)

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Symbol First Second Third

H

He

Li

BeB

C

N

O

F

Ne

1312

2731

520

900800

1086

1402

1314

1681

2080

5247

7297

17572430

2352

2857

3391

3375

3963

11810

148403569

4619

4577

5301

6045

6276

Ionization Energy (IE)

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What factors determine IE Thegreaterthenuclear charge, thegreaterIE.

Greater distance fromnucleusdecreasesIE

Filled and half-filled orbitals have lower energy,

so achieving them is easier, lower IE.

Shielding effect

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Electron affinity is the negative of the energy changethat occurs when an electron is accepted by an atom

in the gaseous state to form an anion.

X (g) + e- X-(g)

F (g) + e- F-(g)

O (g) + e- O-(g)

H = -328 kJ/mol EA = +328 kJ/mol

H = -141 kJ/mol EA = +141 kJ/mol

D. Electron affinity

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Electron Affinity

Definition:The energy given off when aneutral atomin the gas

phase gains an extraelectron to formanegatively charged ion.

1)down a group, electron affinity decreases.

2)across a period, electron affinity increases.

HLi

Na

K

Li Be B C N O F

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Exceptions on electron affinity trends Nonmetals elementsin the first period have lower electron

affinities than the elements below themin their respective groups.

Elements with electron configurations ofXs2, Xp3, and Xp6have

electron affinitiesless than zerobecause they are unusuallystable. e.g. Be, N, Ne

WHY?- Electron affinities are all much smaller than ionizationenergies. Xs2< 0: Stable, diamagnetic atomwith no unpaired electrons.

Xp3< 0:Stable atomwith 3 unpaired p-orbital electrons each occupyingitsown subshell.

Xp6< 0:Stable atomwith filled valence (outermost) shell.

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Group 7-the highest affinity.Group 8-the lowest (zero or ve) valueY???? affinity of Group

2A< 1A and 5A


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Group 3A Elements (ns2np1, n 2)

4Al(s) + 3O2(g) 2Al2O3(s)

2Al(s) + 6H+

(aq) 2Al3+

(aq) + 3H2(g)

Metalloid

Metals

Thus, no reaction with H2O and O2

(HCl)

unipositive

>

stable

than

tripositive

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Sn(s) + 2H+

(aq) Sn2+

(aq) + H2 (g)

Pb(s) + 2H+

(aq) Pb2+

(aq) + H2 (g)

Metalloid

Non-metal

MetalNo reaction

with H2O

(EXAMPLE: HCl)

Form +2 and +4 oxidation state

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N2O5(s) + H2O(l) 2HNO3(aq)

P4O10(s) + 6H2O(l) 4H3PO4(aq)

Non-metal

Metalloid

Metal

N2 forms oxides NO, N2O, NO2, N2O4, N2O5;only N2O5 is solid.

Less reactive metal

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SO3(g) + H2O(l) H2SO4(aq)

Non-metal

Metalloid

Important compounds of S SO2, SO3, H2S

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X + 1e- X-1

X2(g) + H2(g) 2HX(g)

Increasing

reactivity

Nonmetal Formula X2Show high ionization energy and affinity

Anions Halides

(Hydrogen halide)

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-Completely filled outer ns and np subshells(high stability). Highest ionization energyof all elements

-No tendency to accept extra electrons

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Properties of Oxides Across a Period

basic acidicAmphoteric

(both acidic & basic)

5152

Na2O(s) + H2O(l) 2NaOH(aq)

MgO(s) + HCl(aq) Mg2Cl2(aq) + H2O(l)

Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l)basic propertiesAl2O3(s) + 2NaOH(aq)+ 3H2O(l) 2NaAl(OH)4(aq)..acidic

properties

SiO2(s) + 2NaOH(aq) Na2SiO3(aq) + H2O(l)..acidic

Other three oxides are acidic. React with H2OExample:P4O10(s) + 6H2O 4H3PO4(aq)

Conclusion:Moving left to right the period--- - Metallic element decreases

- Basic-Amphoteric- Acidic

_the chemistry of element

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