The Chemical Bond

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The Chemical Bond
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What is a chemical bond?

• It is a force of attraction between atoms, ions
  or molecules.

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How are bonds formed ?

• By electron loss, gain or sharing.

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Which elements lose electrons?

• Metals
• Cations are formed

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Why metals lose electrons ?

• They usually have one, two or three electrons.
• By losing these electrons they can achieve a full
  outer shell and become more stable.
• It is in energy terms easier to lose electrons
  than gain electrons

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Which elements gain electrons?

• Non-metals
• Anions are formed

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Why non-metals gain electrons?

• They usually have five, six or seven electrons.
• By gaining three, two or one electron, they can
  achieve a full outer shell and become more
  stable.
• In energy terms it is easier to gain these
  electrons than lose the outer electrons.

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Bonds

• Covalent
• Polar covalent
• Dative covalent
• Giant covalent
• Ionic
• Metallic

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How are covalent bonds formed?

• By sharing electrons between non-metals.

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Dot and cross diagram-Covalent bond

• Each H atom contributes one electron to the
  shared pair.
• Each H atom has two electrons in its outer shell.
• Each H atom now has a full outer shell

H H
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What is electronegativity?

• It is the ability of an atom to attract the
  bonding electrons in a bond.

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Electronegativity values

• Hydrogen 2.1 Lithium 1.0
• Beryllium 1.5 Boron 2.0
• Carbon 2.5 Nitrogen 3.0
• Oxygen 3.5 Fluorine 4.0

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Which element is most electronegative?

• Fluorine
• The electronegativity is 4.0

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How does the presence of a more electronegative
atom affect a covalent bond?

• The more electronegative atom attracts the shared
  electrons more towards itself.
• The more electronegative atom acquires a partial
  negative charge.
• The less electronegative atom a partial positive
  charge.
• The bond becomes a polar covalent bond.

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Calculate the electronegativity difference for
the following

• C H . Li - F
• O H
• H F
• N H
• C N
• C O
• C - C

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Which bond has the highest electronegativity
difference?

• Li F
• Electronegativity differences greater than or
  equal 1.6 show that the compound is ionic
• This compound is ionic!

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Ionic or polar covalent ?

• This depends on the electronegativity difference
• When the difference is relatively small, the
  bonding is polar covalent.
• When the difference is large, the bonding becomes
  ionic. (1.6 or more)

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How are ionic bonds formed ?

• By exchanging electrons.
• Metals lose electrons and become positive ions.
• Non-metals gain electrons and become negative
  ions.
• The oppositely charged ions attract each other.
  This is ionic bond.

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Ionic bonding
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Ionic bonding
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Giant Covalent Bond

• Structure of diamond

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Giant Covalent Bond

• Structure of diamond

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Giant Covalent Bond

• Structure of Graphite

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How are metallic bonds formed?

• In metals , the outer electrons are free.
• These delocalised electrons form a sea of
  electrons.
• The metal atoms exist as positive ions immersed
  in this sea of electrons.
• The attraction between the positive ions and the
  sea of electrons is known as the metallic bond.

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Metallic bonding

• Metallic bonding in potassium the outermost
  electrons are free and form a sea of
  electrons which attracts the K ions.

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Intermolecular Bonds

• . Van der Waals forces
• . Dipole dipole attraction
• . Hydrogen bonds

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Van der Waals Forces

• Random electron movements creates dipoles.
• Dipoles induces dipoles.
• Process is repeated.
• Oppositely charged dipoles attract.

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Van der Waals Forces
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Van der Waals Forces- Factors

• More electrons, higher forces
• Larger the area of contact, higher forces
• Linear molecules, higher forces
• Branched chains, lower forces
• Spherical molecules, lower forces

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Dipole-dipole attraction

• . More electronegative atom in a covalent bond.
• . Permanent dipoles.
• . Attraction between dipoles.

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Dipole-dipole attraction-Factors

• More the number of electronegative atoms, higher
  attraction
• More electronegative atoms, higher attraction

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Hydrogen Bond

• . Hydrogen in a covalent bond with nitrogen,
  oxygen or fluorine.
• . Permanent dipoles.
• . Attraction between dipoles.
• . Strongest intermolecular force.
• . Intermolecular forces much weaker than normal
  covalent bonds

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Hydrogen bonding in water
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H-bond in H2O

• But for its H-bonding water would be a gas at
  room temperature.
• Due to H-bonding ice floats on water.
• Water freezes with a lot of space in between due
  to H-bonding. This makes ice lighter.

H
O H
H
O
H
H
O
H
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Hydrogen bonding in water
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Hydrogen bonding ..

• Ammonia can form H-bonds.
• HF can also form H-bonds.

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Trend in the melting and boiling point of Group
IV hydrides
mp

• As the elements get larger, there are more
  electrons.
• Van der Waals force is larger.
• MP and BP is larger from CH4 to PbH4.

CH4
PbH4
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Anomalous melting and boiling points of hydrides

• Why ammonia, the lightest hydride has the highest
  melting and boiling point than the corresponding
  hydrides in Group V?
• Why water, the lightest hydride in GroupVI has
  the highest melting and boiling points?
• Why hydrogen fluoride, the lightest hydride in
  Group VII has the highest melting and boiling
  points?

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What are these hydrides

• Gr V NH3, PH3, AsH3, SbH3,BiH3
• Gr VI H2O, H2S, H2Se, H2Te, H2 Po
• Gr VII HF, HCl, HBr, HI, HAt

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Why the mp or bp of NH3 is unusually so high?

• Nitrogen is electronegative.
• NH bonds are polar
• H-bonds formed between ammonia molecules.
• No H-bonds in other molecules, only Van der Waals
  Forces.Partially covalent ionic bonds
• . High charge density of positive ions.
• . Positive ions polarising negative ions.
• . Resulting electron sharing partial
  covalency.Sigma and pi-bonds
• . Sigma bond direct overlap of orbitals
• . Pi- bond sideways overlap of orbitals.

N-H N -H H-bond
BP
BiH3
NH3
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What are dative covalent bonds?

• In a dative covalent bond, both the shared
  electrons come from the same atom.
• The atom giving the electron pair is the donor.
• The atom receiving the electron pair is the
  acceptor.
• There is no difference in length and strength
  between a normal and a dative covalent bond.

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The End