Chapter 6
Ionic Bonds & Some Main-Group Chemistry
Chapter 6
1
Valence Electrons
• When an atom undergoes a chemical
reaction, only the outermost electrons
are involved.
• These electrons are of the highest
energy and are furthest away from the
nucleus. These are the valence
electrons.
• For the main group elements, the
valence electrons are the s and p
electrons beyond the noble gas core.
Chapter 6
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1
Predicting Valence Electrons
• For the main group elements,
the Group number indicates
the number of valence
electrons.
Chapter 6
3
Ion Electron Configurations
• When we write the electron configuration of a cation,
we remove one electron for each positive charge:
Na → Na+
1s2 2s2 2p6 3s1 → 1s2 2s2 2p6
• When we write the electron configuration of an anion,
we add one electron for each negative charge:
O → O21s2 2s2 2p4 → 1s2 2s2 2p6
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2
Ion Electron Configuration: Transition Metals
• Transition metals also lose valence electrons when
forming cations
• However, these elements lose their valence-shell selectrons before losing their d-electrons.
• Electrons with the highest n-quantum number are lost
first.
Fe → Fe2+
[Ar] 4s2 3d6 → [Ar] 3d6
Never
2
[Ar] 4s 3d6 → [Ar] 4s2 3d4
Chapter 6
5
Isoelectronic Ions
• Isoelectronic atoms and ions are those which
have the same number of electrons
• For example:
– A Sodium ion has 10 electrons:
Na → Na+
1s2 2s2 2p6 3s1 → 1s2 2s2 2p6
– A Fluorine ion also has 10 electrons:
F → F1s2 2s2 2p5 → 1s2 2s2 2p6
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3
Ions and Ionic Radii
• Cations are smaller than their parent atom.
– The size decreases because of a decrease in valence shell size,
causing an increase in Zeff
• Anions are larger than their parent atom.
– The size increase is due to an increase the Zeff and in electronelectron repulsion
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7
Ionization Energy
• The ionization energy (EI) of an atom is the amount of energy required to
remove an electron in the gaseous state (units of kJ/mol).
• The closer the electron is to the nucleus, the harder it is to remove.
Core electrons
shield valence
electrons effectively
Electrons in the
same shell shield
one another less
effectively
Chapter 6
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4
Ionization Energy
• There are some
irregularities in the EI
trend
• Why do you think that
Be has a higher EI than
Boron?
• There are additional Ionization Energies,
depending on the number of the electron
you are trying to remove.
• The values for these ionization energies
depend greatly on the location of the
electron in the orbitals, so use your
configurations!
Chapter 6
9
Electron Affinity
• The electron affinity (Eea) of an atom is the amount of energy released
when an electron is added to an atom in the gaseous state (units of
kJ/mol).
• Because the energy is released, electron affinities have negative values.
– The more negative the value, the higher the electron affinity!
The value of an atom’s Eea is
due to a balance between:
1) the attraction between the
new electron and the
nucleus
2) the increase in electronelectron repulsions
Chapter 6
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5
Ionic Bonds and Solids
• An ionic bond is formed when an
element with a small EI value comes
in contact with an element with a
negative Eea value.
• The oppositely charged ions that are
formed by this interaction are then
attracted to one another by
electrostatic forces and are joined by
an ionic bond.
– This “electrostatic force” is similar to
the attraction between opposite poles
on two magnets.
• Ionic bonds are very strong and result
in the formation of a rigid, crystalline
structure called an ionic solid.
Chapter 6
11
Ionic Compounds
• The bonds in an ionic
compounds are very strong
which is why most ionic
compounds are solids at
room temperature
• These compounds have very
high melting points
• In their solid state, these
compounds are not good
conductors of electricity
• However, when the
substance is melted or
dissolved, the crystal is
destroyed, allowing the ions
to roam free and conduct
electricity.
Chapter 6
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6
Lattice Energy (U)
• The lattice energy (U) is the sum of the electrostatic interaction energies
between the ions in a crystal.
• The LE is listed as the amount of energy needed to breakup a crystal and has a
positive value
• If you want to know the amount of energy given off when a lattice forms, just
change the sign to negative!
• The lattice energy is dependent on two factors:
– The charges on the ions (z1 and z2)
– The distance between the ions (d)
U = k•
z1 • z2
d
Chapter 6
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Lattice Energy (U)
One of the below pictures represents NaCl and the other is
MgO. Which is which?
Which of the two crystals do you think has the higher
lattice energy?
Chapter 6
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7
Energetics of Ionic Solid Formation
• For a reaction to proceed, the overall net energy change (∆E)
must be negative
– Merely looking at the EI and Eea of two particular elements is not
enough. Too simple.
For Example: NaCl EI for Na = +495.8 kJ/mo
Eea for Cl = -348.6 kJ/mol
∆E = +147.2 kJ/mol (No Go!)
• The total energetics of ionic reactions can be viewed on a
Born–Haber Cycle which shows how each step contributes
to the overall reaction energy.
Chapter 6
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Energetics of Ionic Solid Formation
Eea
EI
Lattice Energy (U)
Chapter 6
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8
Energetics of Ionic Solid Formation
Chapter 6
17
Energetics of Ionic Solid Formation
• Calculate the overall energy change (in kJ/mol) for the
formation of CaCl from its elements and the formation
of CaCl2 from its elements.
Ea for Cl = -348.6 kJ/mol
EI1 for Ca = +589.8 kJ/mol
EI2 for Ca = +1145 kJ/mol
Heat of Sublimation for Ca = +178.2 kJ/mol
Bond Dissociation Energy for Cl2 = +243 kJ/mol
Lattice Energy for CaCl2 = +2258 kJ/mol
Lattice Energy for CaCl = +717 kJ/mol
• Which is more likely to form, CaCl or CaCl2?
Chapter 6
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9
Chemical Families
• Elements within a group (a column) have similar chemical
properties
• Several of these groups form chemical families
• Trends among these families are most obvious for the main group
elements.
Chapter 6
19
Group 1A: Alkali Metals
• With the exception of hydrogen, all the elements in group 1A are known as the
Alkali metals
• These elements all have one valence s electron, and like to form +1 ions
– This makes them excellent reducing agents
• As you move down the group, the reactivity of the metal increases
• Alkali metals do not look like what we think of as a metal (shiny) because they
react readily with oxygen which coats their outer layer.
• These elements are good conductors of electricity and can be readily formed
into foils and wires.
Chapter 6
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10
Group 1A: Alkali Metals
•
•
Reaction with Halogens
2 M(s) + X2 → 2 MX(s)
Reaction with Oxygen
Forms oxide (Li2O), peroxide (Na2O2), or
superoxide (KO2)
•
Reaction with Hydrogen
2 M(s) + H2 → 2 MH(s)
•
Reaction with Nitrogen
6 Li(s) + N2 → 2 Li3N(s)
Chapter 6
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Group 2A: Alkali Metals
• The elements in group 2A are known as the Alkaline Earth metals
• These elements all have two valence s electrons, and like to form +2 ions
– This makes them excellent reducing agents
• As you move down the group, the reactivity of the metal increases
• These elements are good conductors of electricity and can be readily formed
into foils and wires.
• Compared to Group 1A metals, these metals are smaller, have higher melting
and boiling points and are less reactive.
Chapter 6
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11
Group 2A: Alkali Metals
• Reaction with Halogens
• M(s) + X2 → MX2(s)
• Reaction with Oxygen
• M(s) + O2 → MO(s)
• Reaction with Hydrogen
• 2 Ca(s) + H2 → 2 CaH2(s)
• Reaction with Water (only Ba is vigorous)
• Ba(s) + H2O → Ba2+(aq) + 2 OH–(aq) + H2(s)
Chapter 6
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Group 3A: Boron
•
Elements in Group 3A contain one semimetal that acts as a nonmetal, and four that are
primarily metallic.
•
Boron is acts as a nonmetal and forms covalent bonds.
– It has many similarities to carbon and silicon.
– Its first I.E. is too high (801 kJ/mol) for electron loss so it will not form an ion
– Boron forms covalent compounds not ionic.
•
Always trivalent and never monovalent.
•
Boron can have an incomplete octet (does not follow the octet rule!)
•
Elemental boron is a good semiconductor.
Chapter 6
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12
Group 3A: Aluminum
• Alumnium is the most abundant metal in the earth’s crust
at 8.3%
• Similar to Group 1A and 2A metals, Al is a reducing
agent that undergoes redox reactions by losing all three of
its valence electrons ( 1 p and 2 s electrons) to yield Al3+
ions.
• Aluminum is less reactive than the Group 1A and 2A
metals
• Will react vigorously with O2: Main oxide is Al2O3
• Will react vigorously with halogens: Main halides is AlX3
but can be more complex
Chapter 6
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Group 7A: The Halogens
• The elements in group 7A are known as the Halogens
• These elements are reactive and toxic non-metals
• The Halogens have seven valence electrons, and like to form -1 ions
– This makes them excellent oxidizing agents
• As you move down the group, the reactivity of the non-metal decreases
• These elements exist as solids (I), liquids (Br) or gases (F and Cl) in their
native state.
• The density, melting point and boiling point all increase as you move down
the group.
Chapter 6
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13
Group 7A: The Halogens
Metal Halides:
• Halogens react with every metal in the periodic table to form
metal halides.
2 M + n X2 → 2 MXn
Hydrogen Halides:
• They are formed from salt and sulfuric acid:
2 NaCl(s) + H2SO4(aq) → 2 HCl(g) + Na2SO4(aq)
•
HBr is usually formed from PBr3 and H2O:
PBr3(l) + 3 H2O(l) → 3 HBr(g) + H3PO3(aq)
Chapter 6
Reactivity decreases down the group!
27
Group 8A: The Nobel Gases
•
The elements in group 8A are known as the Nobel Gases
•
These elements are non-metals and have eight valence electrons
•
Because they have eight valence electrons, these elements do not like to undergo redox reactions
at all.
•
These elements exist as gases in their native state.
•
The density, melting point and boiling point all increase as you move down the group.
Reactivity
•
He, Ne and Ar undergo no chemical reactions and form no known compounds.
•
Kr and Xe react only with Fluorine.
–
For example, XeF2, XeF4 and XeF6 are all very powerful oxidizing agents.
Chapter 6
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14
Octet Rule
• The octet rule states that main group elements tend to
undergo reactions that leave them with eight outer
shell electrons.
Chapter 6
29
Octet Rule Exceptions
• There are some exceptions to the octet rule.
– Group 3A non-metals (Boron) can be electron deficient
– Elements in Period 3 or higher and Group 3A or higher can
be electron rich
Chapter 6
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15