Chemical Bonding
Chemical Bonding
• Chemical bonding refers to the
attractive forces that hold atoms
together in compounds.
There are two major classes of
bonding:
• Ionic bonding results from electrostatic
interactions among ions, which often results
from the net transfer of one or more electrons
from one atom to another
-Atoms that lose electrons form positive
ions called cations.
-Atoms that gain electrons form negative
ions called anions.
There are two major classes of
bonding (continued):
• Covalent bonding results from sharing one
or more electron pairs between two atoms
-Atoms that share electron pairs equally
are referred to as nonpolar covalent
bonds.
-Atoms that share electron pairs unequally
are referred to as polar covalent bonds.
All bonds between atoms of different elements
have at least some degree of both ionic and
covalent character.
• Compounds containing predominantly ionic bonding are
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called ionic compounds.
Compounds that are held together mainly by covalent
bonds are called covalent compounds.
Polyatomic ions are groups of covalently bonded atoms
that together form ionic bonds with other atoms.
Some nonmetallic elements, such as H2, N2, O2, F2, Cl2,
Br2, I2 also involve covalent bonding.
Many molecules found in nature consist of covalently
bonded carbon-carbon or carbon-hydrogen bonds or
both. These compounds are classified as organic
compounds.
Properties of Ionic and Covalent
Compounds
Ionic Compounds
Covalent Compounds
• Crystalline solids
• High melting points (>400oC)
• Soluble in polar solvents such as
• Gases, Liquids, or soft solids
• Low melting points (<300oC)
• Insoluble in polar solvents such as
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water
Molten compounds conduct
electricity
Aqueous solutions conduct
electricity
Usually formed between a metal
and a nonmetal.
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water
Most are soluble in nonpolar
solvents such as hexane and
carbon tetrachloride
Liquid, molten, and aqueous
solutions do not conduct electricity
Usually formed between
nonmetals
Representing Compounds
There are several methods used in representing
compounds:
• Chemical formula -gives the number of atoms of
each type in the compound
-Molecular formula gives the actual number of
atoms of each type
-Empirical formula gives the simplest ratio of
atoms of each type
• Structural formula -shows the order in which the
atoms in a compound are connected
Naming Simple Compounds
• Binary Ionic Compounds (Type I)
1. The cation is always named first and
the anion second.
2. A monatomic cation takes its name
from the name of the element.
3. A monatomic anion is named by taking
the root of the element name and adding
–ide.
What is the name of AlCl3?
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1. Aluminum trichloride
2. Aluminum chloride
3. Aluminum chlorine
4. Aluminide chloride
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Naming Simple Compounds
• Binary Ionic Compounds (Type II)- Metals that
can form more than one positive ion can form
more than one type of compound with a given
anion.
1. The charge on the metal ion must be
specified by using a Roman numeral following
the cation name.
2. The anion is named the same as with the
Type I compounds.
What is the name of CuCl?
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1. Copper chloride
2. Copper (I) chloride
3. Copper (II) chloride
4. Copper chlorine
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What is the name of HgO?
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1. Mercury oxide
2. Mercury (I) oxide
3. Mercury (II) oxide
4. Mercury oxygen
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What is the name of Fe2O3?
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1. Iron oxide
2. Iron (I) oxide
3. Iron (II) oxide
4. Iron (III) oxide
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What is the name of MnO2?
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1. Manganese oxide
2. Manganese (I) oxide
3. Manganese (II) oxide
4. Manganese (IV) oxide
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What is the name of GaI3?
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1. Gallium iodide
2. Gallium (I) iodide
3. Gallium (II) iodide
4. Gallium (III) iodide
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Naming Simple Compounds
• Ionic Compounds with Polyatomic Ions
1. Polyatomic ions have specific names
which must be memorized.
2. The other ion is named in the same
manner as the Type I or Type II
compounds.
What is the name of Na2SO4?
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1. Sodium sulfide
2. Sodium sulfate
3. Sodium sulfite
4. Sodium (II) sulfate
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What is the name of Fe(NO3)3?
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1. Iron nitrate
2. Iron (III) nitrate
3. Iron nitrite
4. Iron (III) nitrite
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What is the name of CuCO3?
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1. Copper carbonate
2. Copper (I) carbonate
3. Copper (II) carbonate
4. Copper (III) carbonate
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What is the name of NaOCl?
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1. Sodium oxygen chloride
2. Sodium hypochlorite
3. Sodium chlorite
4. Sodium chlorate
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Naming Simple Compounds
• Binary Covalent Compounds (Type III)
1. The first element in the formula is
named first, using the full element name.
2. The second element is named as if it
were an anion.
3. Prefixes are used to denote the
numbers of atoms present.
Binary Covalent Compounds
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Prefixes
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
Number Indicated
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What is the name of PCl5?
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1. Phosphorus chloride
2. Phosphorus (V) chloride
3. Monophosphorus pentachloride
4. Phosphorus pentachloride
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What is the name of N2O3?
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1. Nitrogen oxide
2. Nitrogen (II) oxide
3. Nitrogen trioxide
4. Dinitrogen trioxide
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What is the name of Nb2O5?
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1. Diniobium pentaoxide
2. Niobium oxide
3. Niobium (II) oxide
4. Niobium (V) oxide
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What is the name of Ti(NO3)4?
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1. Titanium nitrate
2. Titanium tetranitrate
3. Titanium (IV) nitrate
4. Titanium (XII) nitrate
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Writing Formulas
• Given the name of a compound, the following
rules are followed when writing the formula:
1. For ionic compounds, the sum of the charges
must equal zero. Subscripts are added as
needed in order to follow this rule.
2. For covalent compounds, subscripts are
added according to the numerical prefixes given.
Energy of Chemical Bonds
• Bond energy- the energy required to break a
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bond.
Bond energy can be used to indicate the
strength of a bonding interaction.
Ionic bonds form when an atom loses electrons
relatively easily to an atom that has a high
electron affinity.
Ionic compounds result when a metal reacts
with a nonmetal.
Coulomb’s Law
• Coulomb’s law can be used to calculate the energy of
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interaction between a pair of ions.
E = 2.31 x 10-19 J . nm (Q1Q2/ r)
E is energy in joules, r is the radius between the ion
centers in nanometers, and Q1 and Q2 are the numerical
ion charges.
Example: the distance between Na+ and Cl- ions is
0.276 nm. What is the ionic energy per pair of ions?
-8.37 x 10-19 J
A negative sign indicates an attractive force. (The ion
pair has lower energy than the separated ions and is
therefore more stable).
Coulomb’s law can also be used to calculate the
repulsive force between two like charged ions. In this
case, the sign will be positive.
Energy Effects in Binary Ionic
Compounds
• Lattice energy is the change in energy
that takes place when separated gaseous
ions are packed together to form an ionic
solid. (Energy that is released when an
ionic solid forms from its ions).
• If energy is released, it represents an
exothermic process. (Energy will have a
negative sign).
Lattice Energy Calculations
• A modified form of Coulomb’s Law can be
used to calculate lattice energy.
• Lattice energy = k(Q1Q2 / r)
• K is a constant that depends on the
structure of the solid and electron
configuration of the ions, Q1 and Q2 are
the charges on ions, and r is the distance
between the nuclei of the ions.
Which compound has the most
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exothermic lattice energy?
1. NaCl
2. KCl
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Which compound has the most
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exothermic lattice energy?
1. LiF
2. LiCl
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Which compound has the most
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exothermic lattice energy?
1. Mg(OH)2
2. MgO
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Which compound has the most
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exothermic lattice energy?
1. Fe(OH)2
2. Fe(OH)3
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Which compound has the most
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exothermic lattice energy?
1. NaCl
2. Na2O
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Which compound has the most
exothermic lattice energy?
1. MgO
2. BaS
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Which compound has the most
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exothermic lattice energy?
1. MgO
2. NaF
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Covalent Bonding
• Electrons are shared between the nuclei of two atoms.
• Polar covalent bonding occurs when there is unequal
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sharing between the atoms.
Unequal sharing results from electronegativity
differences between atoms.
Electronegativity is the ability of an atom in a molecule
to attract shared electrons to itself.
Electronegativity increases bottom to top and left to
right across the periodic table.
The fractional charge that results is called a dipole
moment.
The dipolar character of a molecule is represented by an
arrow pointing to the negative charge center.
See page 334 for table of electronegativities.
Which of the following pairs of elements
forms a nonpolar covalent bond?
1. H-H
2. S-H
3. O-H
4. F-H
5. Cl-H
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Select the answer choice that lists the pairs of
elements in order of increasing polarity.
1. H-H, O-H, Cl-H, S-H, F-H
2. F-H, Cl-H, O-H, S-H, H-H
3. H-H, O-H, S-H, Cl-H, F-H
4. H-H, S-H, Cl-H, O-H, F-H
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Polar and Nonpolar Molecules
• A molecule is nonpolar if the dipoles
cancel (opposed bond polarities cancel
out).
• This occurs when the bonds are
symmetrically arranged within the
molecule.
• A molecule is polar if the dipoles do not
cancel out.
Which of the following
molecules is nonpolar?
1. H2O
2. CH4
3. CH3Cl
4. HCl
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Which of the following
molecules is polar?
1. CO2
2. CO
3. C6H6
4. O2
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Types of Covalent Bonds
• Single bond-One pair of electrons is
shared between atoms
• Double bond-two pairs of electrons are
shared between atoms
• Triple bond-three pairs of electrons are
shared between atoms
Bond Energy and Bond Length
• Bond energy increases as follows:
single bond < double bond < triple bond
• Bond length increases as follows:
triple bond < double bond < single bond
Bonding Models
• Localized Electron Bonding Model
•
Molecules consist of lone pairs and bonding pairs
of electrons
VSEPR Model
The structure around a given atom is
determined principally by minimizing electron
pair repulsions. (Bonding and nonbonding pairs
are positioned as far apart as possible)
Lewis Structures
• The Lewis structure shows how the
valence electrons are arranged among the
atoms in a molecule.
• For the formation of a stable compound,
the atoms should achieve noble gas
configuration.
Steps for Writing Lewis
Structures
• Find the total number of electrons needed if each
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element had a complete octet. (remember hydrogen
only needs two valence electrons to be stable)
Determine the total number of valence electrons
available.
Find the difference between what is needed and what is
available in order to determine the number of electrons
shared.
Every two electrons shared represents one bond.
Give the Lewis structure for
each of the following.
• H2O
• CH4
• HCl
• NO+
• NH3
• HCN
• CO2
Exceptions to the Octet Rule
• Some compounds form in which one or more
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atoms will have fewer than 8 electrons in the
outer level (incomplete octet)
Example: BF3
Other compounds form in which one atom may
have more than 8 electrons in the outer level
(Exceeds the octet rule)
Example: SF6
Resonance
• Resonance is invoked when more than one valid
Lewis structure can be written for a molecule.
• The resulting electron structure of the molecule is
given by an average of these resonance structures.
• This is supported by experimental observations in
which the bond lengths and strengths are between
those expected if the bonds were single and/or
double.
Formal Charge
• Molecules that exceed the octet rule can
often have many nonequivalent Lewis
structures, all which obey the rules for
writing Lewis structures.
• Through the determination of formal
charges, the Lewis structures can be
evaluated as to which best describes the
bonding in the molecule.
Rules Governing Formal Charge
• Determine the sum of the lone pair electrons and one•
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half of the shared electrons. This is the number of
valence electrons assigned to the atom.
Subtract the number of valence electrons assigned to
the atom from the valence electrons on the free , neutral
atom.
The sum of the formal charges must equal the overall
charge on the ion or molecule.
If nonequivalent structures exist, those with formal
charges closest to zero and with any negative formal
charges on the most electronegative atoms are
considered the best description of the bonding.
Sigma bonds
Pi bonds
Double bond
Triple
Bond