TYPES OF CHEMICAL BONDS

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TYPES OF CHEMICAL BONDS

• IONIC BONDS
• COVALENT BONDS
• HYDROGEN BONDS
• METALLIC BONDS

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IONIC BONDING
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IONIC BONDING
When an atom of a nonmetal takes one or more
electrons from an atom of a metal so both atoms
end up with eight valence electrons
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IONIC BONDING
IS THE COMPOUND AN IONIC COMPOUND?
Mg N
3
2
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IONIC BOND FORMATION
Non-Metal
Metal
Neutral atoms come near each other. Electron(s)
are transferred from the Metal atom to the
Non-metal atom. They stick together because of
electrostatic forces, like magnets.
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IONIC BONDING

• Metals will tend to lose electrons and become
• POSITIVE CATIONS

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IONIC BONDING

• Nonmetals will tend to gain electrons and become
• NEGATIVE ANIONS

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Properties of Ionic Compounds

• Crystalline structure.
• A regular repeating arrangement of ions in the
  solid.
• Ions are strongly bonded.
• Structure is rigid.
• High melting points- because of strong forces
  between ions.

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Crystalline structure
The POSITIVE CATIONS stick to the NEGATIVE
ANIONS, like a magnet.
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Do they Conduct?

• Conducting electricity is allowing charges to
  move.
• In a solid, the ions are locked in place.
• Ionic solids are insulators.
• When melted, the ions can move around.
• Melted ionic compounds conduct.
• Melting points always above 800ºC.
• Dissolved in water they conduct.

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Ionic solids are brittle
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Ionic solids are brittle

• Strong Repulsion breaks crystal apart.

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COVALENT BONDING
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COVALENT BONDING
When an atom of one nonmetal shares one or more
electrons with an atom of another nonmetal so
both atoms end up with eight valence electrons
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COVALENT BOND FORMATION
When one nonmetal shares one or more electrons
with an atom of another nonmetal so both atoms
end up with eight valence electrons
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COVALENT BONDING
IS THE COMPOUND A COVALENT COMPOUND?
C O
2
YES since it is made of only nonmetal elements
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Covalent bonding

• Fluorine has seven valence electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• Both end with full orbitals

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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• Both end with full orbitals

F
F
8 Valence electrons
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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• Both end with full orbitals

F
F
8 Valence electrons
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Single Covalent Bond

• A sharing of two valence electrons.
• Only nonmetals and Hydrogen.
• Different from an ionic bond because they
  actually form molecules.
• Two specific atoms are joined.
• In an ionic solid you cant tell which atom the
  electrons moved from or to.

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Water

• Each hydrogen has 1 valence electron
• Each hydrogen wants 1 more
• The oxygen has 6 valence electrons
• The oxygen wants 2 more
• They share to make each other happy

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Water

• Put the pieces together
• The first hydrogen is happy
• The oxygen still wants one more

H
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Water

• The second hydrogen attaches
• Every atom has full energy levels

H
H
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Multiple Bonds

• Sometimes atoms share more than one pair of
  valence electrons.
• A double bond is when atoms share two pair (4) of
  electrons.
• A triple bond is when atoms share three pair (6)
  of electrons.

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Carbon dioxide

• CO2 - Carbon is central atom
• Carbon has 4 valence electrons
• Wants 4 more
• Oxygen has 6 valence electrons
• Wants 2 more

C
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Carbon dioxide

• Attaching 1 oxygen leaves the oxygen 1 short and
  the carbon 3 short

C
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Carbon dioxide

• Attaching the second oxygen leaves both oxygen 1
  short and the carbon 2 short

C
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Carbon dioxide

• The only solution is to share more

C
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Carbon dioxide

• The only solution is to share more

C
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Carbon dioxide

• The only solution is to share more

C
O
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Carbon dioxide

• The only solution is to share more

C
O
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Carbon dioxide

• The only solution is to share more

C
O
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Carbon dioxide

• The only solution is to share more

C
O
O
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Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom gets to count all the atoms in the bond

C
O
O
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Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom gets to count all the atoms in the bond

8 valence electrons
C
O
O
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Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom gets to count all the atoms in the bond

8 valence electrons
C
O
O
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Carbon dioxide

• The only solution is to share more
• Requires two double bonds
• Each atom gets to count all the atoms in the bond

8 valence electrons
C
O
O
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Coordinate Covalent Bond

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide
• CO

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Coordinate Covalent Bond

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide
• CO

O
C
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Coordinate Covalent Bond

• When one atom donates both electrons in a
  covalent bond.
• Carbon monoxide
• CO

O
C
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Coordinate covalent bond

• Other examples

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Polar Bonds

• When the atoms in a bond are the same, the
  electrons are shared equally.
• This is a nonpolar covalent bond.
• When two different atoms are connected, the atoms
  may not be shared equally.
• This is a polar covalent bond.
• How do we measure how strong the atoms pull on
  electrons?

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Electronegativity

• A measure of how strongly the atoms attract
  electrons in a bond.
• The bigger the electronegativity difference the
  more polar the bond.
• 0.0 - 0.3 Covalent nonpolar
• 0.3 - 1.67 Covalent polar
• gt1.67 Ionic

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How to show a bond is polar

• Isnt a whole charge just a partial charge
  figure it out by drawing LEWIS STRUCTURES
• d means a partially positive (no lone pairs)
• d- means a partially negative (atom has lone
  pairs)
• The Cl pulls harder on the electrons
• The electrons spend more time near the Cl

d
d-
H
Cl
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Polar Molecules

• Molecules with a positive and a negative end
• Requires two things to be true
• The molecule must contain polar bonds
• This can be determined from differences in
  electronegativity.
• Symmetry can not cancel out the effects of the
  polar bonds.
• Must determine geometry first.

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Is it polar?

• HF
• H2O
• NH3
• CCl4
• CO2

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Intermolecular Forces

• What holds molecules to each other

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Intermolecular Forces

• They are what make solid and liquid molecular
  compounds possible.
• The weakest are called van der Waals forces -
  there are two kinds
• Dispersion forces
• Dipole Interactions
• depend on the number of electrons
• more electrons stronger forces
• Bigger molecules

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Dipole interactions

• Depend on the number of electrons
• More electrons stronger forces
• Bigger molecules more electrons
• Fluorine is a gas
• Bromine is a liquid
• Iodine is a solid

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Dipole interactions

• Occur when polar molecules are attracted to each
  other.
• Slightly stronger than dispersion forces.
• Opposites attract but not completely hooked like
  in ionic solids.

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Dipole interactions

• Occur when polar molecules are attracted to each
  other.
• Slightly stronger than dispersion forces.
• Opposites attract but not completely hooked like
  in ionic solids.

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Dipole Interactions
d d-
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Hydrogen bonding

• Are the attractive force caused by hydrogen
  bonded to F, O, or N.
• F, O, and N are very electronegative so it is a
  very strong dipole.
• The hydrogen partially share with the lone pair
  in the molecule next to it.
• The strongest of the intermolecular forces.

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Hydrogen Bonding
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Hydrogen bonding
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Metallic Bonds

• How atoms are held together in the solid.
• Metals hold onto there valence electrons very
  weakly.
• Think of them as positive ions floating in a sea
  of electrons.

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Sea of Electrons

• Valence electrons are free to move through the
  solid.
• Metals conduct electricity.

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Metals are Malleable

• Hammered into shape (bend).
• Ductile - drawn into wires.

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Malleable
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Malleable

• Electrons allow atoms to slide by.