Chapter 8 Covalent Bonding

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Chapter 8Covalent Bonding

• General Chemistry I
• T.ARA

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Chemical Bonding

• Now that we know something about electron
  configurations, we can take a closer look at the
  ways atoms form bonds.
• There are two main types of chemical bonds
• Ionic Bonds
• Covalent Bonds

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A. Ionic Bonding

• Ionic bonding involves the transfer of valence
  electrons from one atom (usually a metal) to
  another atom (a nonmetal) such that each atom
  gains a noble gas configuration.
• The ionic bond is the electrostatic attraction
  between the cation () and the anion (-) that
  result from the electron transfer.
• The two bonded atoms do not share electrons.

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A. Ionic Bonding Sodium Chloride

• A sodium atom will readily lose a valence
  electron
• Na Ne3s1 ? Na Ne e-
• A chlorine atom will readily accept an electron
• Cl Ne3s23p5 e- ? Cl- Ar
• The sodium cation the chloride anion are held
  together by an electrostatic (coulombic)
  attraction opposite charges attract.
• Na Cl- ? NaCl (ionic bond)

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A. Ionic Bonding Sodium Chloride

Using Lewis symbols
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1. Polyatomic Ions

• Polyatomic Ion a group of covalently bonded
  atoms with an overall positive or negative charge
• The atoms within
• a polyatomic ion are
• covalently bonded
• together, but
• polyatomic ions form
• ionic bonds with other
• ions.

eg. NO3- is a polyatomic anion.
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B. Covalent Bonding

• In a covalent bond the electrons are the glue
  that holds the atoms together.
• A covalent bond is formed when two atoms share
  one or more pairs of electrons.
• Each atom will form enough covalent bonds to
  achieve a noble gas configuration.

G.N. Lewis
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B. Covalent Bonding Hydrogen

• H H ? H H
• The two electrons are shared evenly between the
  two hydrogen atoms.
• It is as if each atom has two electrons the
  noble gas configuration of He.
• H H

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B. Covalent Bonding Hydrogen

• The actual electron probability cloud (?2) for
  the two electrons in the H-H bond looks like this.

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1. Lewis Structures

• A Lewis structure is a way of drawing a molecule
  that shows all valence electrons as dots or lines
  that represent covalent bonds.
• eg. The Lewis structure for H2 can be drawn in
  two ways
• H H or HH
• a) A single line represents two covalently shared
  electrons also known as a single bond.

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1. Lewis Structures

• The Lewis Structure for F2
• Two electrons are shared between the two F atoms
  (one single covalent bond).
• Each F atom also has three unshared electron
  pairs. These non-bonding electron pairs are
  called lone pairs.

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2. The Octet Rule

• Note that by sharing electrons, it is as if each
  F atom has eight electrons - the noble gas
  configuration of Ne.
• The Octet Rule Main group elements with more
  than two valence electrons gain, lose, or share
  electrons to achieve a noble gas configuration
  characterized by eight valence electrons.

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Draw a Lewis structure for water (H2O) that obeys
the octet rule.
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Draw a Lewis structure for ammonia (NH3) that
obeys the octet rule.
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Draw a Lewis structure for methane (CH4) that
obeys the octet rule.
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2. The Octet Rule

• As the previous examples illustrated, the number
  of covalent bonds an atom must form to achieve an
  octet is equal to eight minus it group number.

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a) Multiple Covalent Bonds

• It is not always possible for atoms to gain a
  full octet by sharing single electron pairs with
  other atoms.
• In other words, it is not always possible to
  construct a valid Lewis structure using only
  single bonds.
• eg. N2

In this structure, each nitrogen atom would have
only six valence electrons two short of an
octet.
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a) Multiple Covalent Bonds

• Two atoms can share more than one electron pair
  to gain a full octet.
• Double Bond when 2 electron pairs (4 electrons)
  are shared between 2 atoms
• Triple Bond when 3 electron pairs (6 electrons)
  are shared between two atoms
• Double triple bonds are referred to as multiple
  covalent bonds.

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a) Multiple Covalent Bonds

• Ethane (C2H6), Ethylene (C2H4) Acetylene (C2H2)

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a) Multiple Covalent Bonds
4 total bonds
3 total bonds
2 total bonds
1 bond
When necessary, atoms will form any combination
of single, double and triple covalent bonds to
gain a full octet.
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Draw a valid Lewis structure for formaldehyde
(CH2O).
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Draw a valid Lewis structure for the cyanide ion
(CN-).
Lewis Structures of Ions Add one additional
valence electron for every negative charge
subtract one valence electron for every positive
charge.
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b) Formal Charge

• As you just saw, even when the octet rule is
  obeyed, some compounds have an overall charge.
• This means that the compound contains one or more
  charged atoms.
• Formal charge the charge a bonding atom would
  have if its bonding electrons were shared equally
• Formal Charge Atomic Group
• lone pair electrons
• ½ ( bonding electrons)

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Draw all possible Lewis structures for N2O
(O-N-N) assign formal charges.
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• Draw all possible Lewis structures for N2O
  (O-N-N) assign formal charges.

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3. Resonance Structures

• When more than one Lewis structure can be drawn
  for a molecule, the structures are called
  resonance structures.
• Each resonance structure contributes to the
  overall structure of the molecule.
• Individual resonance structures DO NOT ACTUALLY
  EXIST we use resonance structures conceptually
  to help us understand molecular structures!!!
• The actual structure is a weighted average of the
  resonance structures called a resonance hybrid.

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3. Resonance Structures
eg. Resonance Structures of Ozone

• Each bond is in between a single a double bond.
• Each terminal O has a partial negative charge.

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3. Resonance Structures

• Resonance structures are connected by
  double-headed arrows ( ).
• Each resonance structure must have the same
  overall charge as the molecule.
• Lower energy resonance structures contribute more
  to the overall structure of the molecule.
• Resonance structures are lowest in energy when
• All atoms with full octets
• The minimum of formal charges
• Negative charges on electronegative atoms (more
  on this in a minute)

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3. Resonance Structures

• Which N2O resonance structure(s) are lowest in
  energy?

A
B
C
B C both have only two formal charges lower
in energy than A.
High in Energy Too many formal charges!!
C is lowest in energy because the negative charge
is on O (more electronegative than N) more on
this in a minute!
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4. Bond Lengths Bond Strengths

• How does the type of bond (bond order) affect the
  properties of the bond?
• Bonds get shorter as the number of electrons
  shared between the two atoms increases. (Bonds
  get shorter as the bond order increases.)
• Bond Length
• Single Bond (1) gt Double Bond (2) gt Triple Bond
  (3)

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4. Bond Lengths Bond Strengths

• The bond order also affects the strength of the
  bond (the bond enthalpy).
• As the bond order increases, the bond gets
  stronger (harder to break).

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C. Ionic vs. Covalent Bond Polarity

• Now we know the difference between ionic and
  covalent bonds.
• Given a specific compound, how do you know which
  type of bond to expect?

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1. The Continuum

• In reality, all bonds have some covalent
  character some ionic character.
• In other words, all bonds fit somewhere on a
  continuum between covalent ionic.
• The real questions are
• How evenly are the electrons being shared between
  the two atoms?
• How can you predict how evenly the electrons will
  be shared?

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1. The Continuum

• Covalent Polar Ionic
• Covalent
• - electrons shared - electrons shared -
  electrons not
• equally unequally shared at
  all
• - No separation - some separation -
  complete separation
• of charge of charge of charge
• - held together by - held together by -
  held together by
• shared electrons shared electrons
  electrostatic
• attraction

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2. Electronegativity

• In 1932, Linus Pauling proposed the idea of
  electronegativity to explain the ways atoms share
  electrons in bonds.

Electronegativity the ability of an atom in a
covalent bond to attract shared electrons to
itself the electron-pulling power of an atom
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2. Electronegativity

The higher the EN, the more tightly an atom holds
its electrons.
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a) Bond Polarity

• The polarity of a bond describes how evenly the
  electrons are shared between the two atoms.
• The more polar a bond, the less evenly the
  electrons are shared.
• A polar bond is indicated by using ? and ?- to
  represent the partial charges on the atoms.

? less electronegative atom ?- more
electronegative atom
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a) Bond Polarity

• The greater the difference in electronegativity
  between the two atoms, the more polar the bond
  between them will be.

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a) Bond Polarity

• Approximate Guidelines
• If ?EN 0.5, the electrons are shared fairly
  evenly the bond is nonpolar covalent.
• If ?EN ? 1.9, the electrons are localized on the
  more electronegative atom the bond is ionic.
• If 1.9 gt ?EN gt 0.5, the electrons are shared
  unequally the bond is polar covalent.

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The Continuum
H2 LiF HCl MgO CsI N2
Covalent
Polar Covalent
Ionic

?EN
0
3.0
0.5
1.0
1.5
2.0
2.5
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D. Exceptions to the Octet Rule

• Most structures containing main group elements
  follow the octet rule, but there are exceptions
  in certain compounds
• Some atoms have fewer than eight electrons (less
  than an octet).
• Some atoms have more than eight electrons (more
  than an octet).

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1. Less than an Octet

• Elements in group 3A have three valence
  electrons.
• Using those three electrons to form three bonds
  gives the central atom only six valence electrons
  (less than an octet).
• eg. BH3

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1. Less than an Octet

• Because boron has less than an octet, BH3 is very
  reactive it will react to form an octet.
• The nitrogen in ammonia uses its lone pair
  electrons to form a bond with boron both B and
  N now have an octet.

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2. More than an Octet Expanded Valence

• Elements in the third period lower can form
  stable compounds in which the central atom has
  more than 8 electrons an expanded valence.
• Unlike elements in the first two periods (like N
  O), third row elements (like P S) can use
  their empty 3d orbitals to accommodate extra
  electrons.

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2. More than an Octet Expanded Valence

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2. More than an Octet Expanded Valence

• eg. PCl5, SBr6, ClF3