_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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Group 4A Elements (ns2np2, n 2)
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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Group 5A Elements (ns2np3, n 2)
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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Group 6A Elements (ns2np4, n 2)
SO3(g) + H2O(l) H2SO4(aq)
Non-metal
Metalloid
Important compounds of S SO2, SO3, H2S
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Group 7A Elements (ns2np5, n 2)
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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Group 8A Elements (ns2np6, n 2)
-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