Transcript CH 7 Periodic Table Properties

Periodic Variation in Physical
Properties
Sizes of atoms and ions
Ionization energy
Electron affinity
Metallic properties
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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
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Effective Nuclear Charge
Effective Nuclear charge, Zeff: the actual nuclear charge
a valence electron experiences.
In a many-electron atom,
the Zeff depends on two
factors:
1. Attraction between
electrons and nucleus
2. Repulsion between
electrons in orbitals
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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.
It depends on:
Size of the atom: Coulomb’s Law
Penetration of orbitals: (10% probability being next to
the nucleus)
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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
•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?
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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.
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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.
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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
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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?
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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
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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 nincreases
– Atoms radii increase in a
group
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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
Rb
Atomic Radius (nm)
K
Na
Ga
Li
V
Zn
Ar
Ne
H
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Atomic Number
Kr
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.
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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.
Ge Si
Se
Cl
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
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Ionic Radius
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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.
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8.3
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8.3
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)
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Radii of Atoms and Corresponding Ions
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.
N3- Li+
C
O2-
1. C < Li+ < O2- < N32. N3- < O2- < C < Li+
3. Li+ < C < N3- < O24. Li+ < C < N3- < O25. Li+ < C < O2- < N3-
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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 shell
Filled 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 sorbital
– 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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Successive Ionization Energies (kJ/mol)
IE1
IE2
IE3
IE4
IE5
IE6
IE7
Na Mg
Al
496 738
578
4562 1451 1817
6912 7733 2745
9543 10540 11575
13353 13630 14830
16610 17995 18376
20114 21703 23293
Si
787
1577
3231
4356
16091
19784
23783
P
1012
1903
2912
4956
6273
22233
25397
S
Cl
Ar
1000 1251 1520
2251 2297 2665
3361 3822 3931
4564 5158 5770
7013 6540 7238
8495 9458 8781
27106 11020 11995
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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?
I = ionization energy
1.
2.
3.
4.
5.
I1 of Na
I2 of Na
I1 of Mg
I2 of Mg
I3 of Mg
#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- → NeΔE = 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 sorbital.
– 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
•
Metalloids
• Nonmetals
– malleable & ductile  Also known as
semi-metals
– shiny, lustrous
– conduct heat and  Show some
metal and some
electricity
nonmetal
– most oxides basic
properties
and ionic
– form cations in
solution
– lose electrons in
reactions oxidized
 brittle in solid state
 dull
 electrical and
thermal insulators
 most oxides are
acidic and molecular
 form anions and
polyatomic anions
 gain 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.
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Properties of Metal, Nonmetals,
and Metalloids
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Metals versus Nonmetals
Differences between metals and nonmetals tend
to revolve around these properties.
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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.
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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.
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Alkali Metals
• Found only as compounds in nature.
• Have low densities and melting points.
• Also have low ionization energies.
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Alkali Metals
Their reactions with water are famously exothermic.
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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.
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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
– O22−, peroxide
– O21−, superoxide
• Tends to take electrons
from other elements
(oxidation)
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Sulfur
• Weaker oxidizing
agent than
oxygen.
• Most stable
allotrope is S8, a
ringed molecule.
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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.
2.
3.
4.
5.
CO2
P 2 O5
BaO
XeO3
SO2
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Periodic Trends Interactive
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