Secs 7.3-8: Periodic Properties
Review: electron configurations and effective nuclear
charge
Goal: relate metallic/nonmetallic character of elements
to position on the Table in terms of electron
configuration and Zeff
We will focus on 3 properties:
Atomic/ionic radii
Ionization energy
Electron affinity
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Atomic and ionic radii
How to accurately determine the radius of an atom?
Use an alternate definition: e.g., for N2: treat the N-N
bond length as the sum of atomic radii of 2 N atoms to
give the bonding atomic radius
What are the periodic trends in bonding atomic radii?
note that:
Li > Be > B > C > N > O > F
Na > Mg > Al > Si > P > S > Cl
(Check out Fig 7.6...)
In general:
within each group: atomic radius increases from top
to bottom
within each period: atomic radius decreases from left
to right
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What factors determine the atomic radius?
what determines the size of the outermost orbital?
Orbital size is determined by:
Principal quantum number
Effective nuclear charge
Increase n: increase orbital size
Increase Zeff: decrease orbital size
Moving across a period: filling orbital with same n
value; what happens to Zeff?
Moving down a group: changing n value of
outermost orbital; what happens to Zeff?
e.g., arrange the following atoms in order of
increasing atomic radius:
Ca, Mg, Be
P, Cl, Sr
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Which has a larger effect on radius: moving one place
across a period or moving down one group? Why?
Sizes of Ions
Very important in determining structure/stability of ionic
compounds
Why? Recall force of attraction between charges:
F
kQ1Q2
d2
How does this affect the attractive force between ions
in a compound?
Ionic size depends on:
nuclear charge
# of e- in the ion
orbitals in which outermost e- reside
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In general,
cations are smaller than their parent atoms
anions are larger than their parent atoms
Periodic trends in ionic size
for ions of the same charge, size increases moving
down a group
isoelectronic series: ions with same # of ee.g., O2-, F-, Na+, Mg2+, Al3+
what is the trend in size in this series? why?
e.g., arrange the following atoms & ions in order of
increasing size:
Ar, Cl-, S2-, K+
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Electron configurations of ions
When a cation or anion is formed, e- are
added/removed from the orbital with the largest value of n:
E.g., write electron configurations for Na+, S2-
Ionization of transition-metal ions
different than representative elements
most transition metal atoms do not ionize to noblegas configurations
transition metal atoms lose their valence shell s efirst, followed by d e-
e.g., write the electron configurations for
Cr3+
Co2+
V2+
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Ionization energies
The first ionization energy I1 is
defined (in the gas phase) by
the reaction
M(g)  M(g)+ + e-
The second ionization energy I2 is defined by
M(g)+  M(g)2+ + elarger value of I: stronger binding of e- to atom or
ion
Removal of an electron always requires energy!
Consider the following ionization energies
(kJ/mol):
element
I1
I2
I3
Na
496
4560
Mg
738
1450
7730
Al
577
1816
2744
Removal of successive e- requires more energy: why?
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What happens to I after a noble-gas configuration is
reached?
Periodic trends in ionization energies
Within each period:
I1 generally increases with atomic number
alkali metals: lowest I.E.
noble gases: highest I.E.
Within each group:
I.E. decreases with increasing atomic number
How to explain these trends?
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I.E. depends on both Zeff and distance from e- to nucleus
moving left to right: Zeff increases, atomic radius decreases
moving down: Zeff constant, atomic radius increases
E.g., Based on position on the periodic table, which of the
following pairs would have the largest first ionization
energy?
Si, C
K, Cr
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Electron affinity
ionization: M(g)  M(g)+ + emeasures energy changes associated with removing
e-
electron affinity, E: M(g) + e-  M-(g)
measures energy changes that occur when e- is added
to atom or ion
Sign convention for electron affinity (EA)
For cations and most neutral atoms: adding e- is
exothermic; EA is generally negative
For anions and some neutral atoms: EA positive
(work must be done to force e- onto atom)
Notice that....
Groups 2A, 8A: E is positive
Group 7A (halogens): E is strongly negative
why?
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Periodic trends in electron affinities
In general, EA becomes more negative as we move
across a period
'exceptions': atoms with filled or half-filled subshells,
e.g., Be, N, Ne
Moving down a group: EA does not change greatly
why?
Periodic properties of groups of atoms
Three broad catagories of elements
Metals
Semimetals
Nonmetals
What are the properties of these groups? What
differentiates them?
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Metals
Properties
Metals tend to have low ionization energies and tend to
form cations easily (easily oxidized)
They tend to be large in size
They tend to have only slightly negative electron
affinities
We will use these three properties - low IE, large size,
slightly negative EA – to define metallic character
What is the periodic trend in metallic character? Where
are the most metallic elements? The least metallic?
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General trends in reactivity of metals
Metals react with nonmetals to form solid ionic
compounds (salts):
M(s) + NM(g, l)  ionic compound(s)
A particularly important reaction of this type is between a
metal and oxygen to form a metal oxide (MO), e.g.
4Fe(s) + 3O2(g)  2Fe2O3(s)
Metal oxides are basic, ionic solids; those that dissolve in
H2O form metal hydroxides (MOH), e.g.,
MO(s) + H2O(l)  MOH(aq)
e.g., write a balanced equation for the reaction of iron (III)
oxide with water.
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Metal oxides react with acids to form salts and H2O:
MO(s) + acid(aq)  salt(aq) + H2O(l)
e.g., write a balanced equation for the reaction of nickel (II)
oxide with aqueous hydrochloric acid
Active metals (e.g. groups 1A and 2A) react with water to
form MOH and hydrogen:
M(s) + H2O(l)  MOH(aq) + H2(g)
E.g., write a balanced equation for the reaction between
solid potassium and water.
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Active metals will also react with aqueous acid to form a
salt and hydrogen gas (cf. activity series)
M(s) + acid(aq)  salt(aq) + H2(g)
E.g., write a balanced equation for the reaction between
solid aluminum and nitric acid.
Summary: Metals are characterized by:
_______ ionization energies
_______ size
_______ electron affinities
Metallic character ____________ moving right to left
across a period and ___________ moving down a group.
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Metals react with nonmetals to form ____________
Metal oxides are _________ oxides.
Metal oxides react with acids to form _______ and
_______.
Active metals react with water to form __________ and
_________.
Active metals react with acids to form ___________ and
_________.
Notice: the more metallic an element, the more basic its
oxide and the more reactive the element will be towards
acid and water!
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Group trends for the active metals
Group 1A: alkali metals:
Lowest I1 of all the elements
Exist in nature only as compounds
Form solid hydrides, sulfides, and chlorides
Form oxides/peroxides/superoxides in reactions
with oxygen
Group 2A: alkaline earth metals
Less reactive than alkali metals (higher I1)
Lighter species (Be, Mg): do not react with H2O(l)
Heavier species (Ca, Sr, Ba) react readily with
H2O(l)
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Nonmetals
Properties
Tend to have very large
ionization energies
Tend to be small in size
Tend to have very
negative electron
affinities – gain electrons
and are easily reduced
We will use these three properties - high IE, small size,
very negative EA – to define nonmetallic character
What is the periodic trend in nonmetallic character?
Where are the most nonmetallic elements? The least
nonmetallic?
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Reactions of nonmetals
tend to form anions in reactions with metals:
Metal(s) + Nonmetal(g,l) salt(s)
Nonmetal oxides (NMO) are acidic oxides: dissolve in H2O
to form acids
NMO(g,s) + H2O(l)  acid(aq)
e.g., write a balanced equation for the reaction of carbon
dioxide with water.
Summary: Nonmetals are characterized by:
_______ ionization energies
_______ size
_______ electron affinities
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Nonmetallic character ____________ moving left to right
across a period and ___________ moving up a group
Nonmetal oxides are ____________ oxides.
Notice: the more nonmetallic an element, the more acidic
its oxide!
E.g., Arrange the following oxides in order of increasing
acidity:
CO2, MgO, Al2O3, SO3, BaO, P2O5
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Group Trends for Nonmetals
Hydrogen
Forms +1 ion with nonmetals
Forms -1 ion with active metals – much higher
ionization energy than metals
Forms solid hydrides with metals
Group 6A:
Oxygen: two forms (or allotropes)
Molecular oxygen, O2
Ozone, O3
Generally forms oxide ion
Also forms peroxide ion O22- and
superoxide ion O2Sulfur
Exists as 8-membered rings; brittle
molecular solid (usually just write S(s)
rather than S8(s))
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Group 7A: halogens
all are typical nonmetals
consist of diatomic molecules: F2(g),
Cl2(g), Br2(l), I2(s)
EA is the most
negative of the
elements; tend
to form halide
anions
Group 8A: Noble gases
Exist as monoatomic gases in nature
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Problems du Jour
Arrange the following atoms in order
of increasing atomic radius:
Mg, F, P, O, Ca
In, Ar, Rb, Ge, He
Why is the second ionization energy of K much
greater than its first ionization energy?
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Problems du Jour
Write balanced equation for the following reactions:
Lithium with water
Calcium oxide with water
Copper (II) oxide with nitric acid
Aluminum with chlorine gas
Sulfur with oxygen
Sulfur trioxide with water
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