CHEMICAL BONDING

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CHEMICAL BONDING

• 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
ION any atom with more or less electrons that
it is supposed to have
Remember that the number of electrons is
supposed to be equal to the number of Protons if
the atom has a neutral charge
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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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IONIC BONDING

• Na1 is called a sodium ion
• The 1 symbol means it
• has lost one electron

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IONIC BONDING

• Mg2 is called a magnesium ion
• The 2 symbol means it
• has lost two electron

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IONIC BONDING

• S-2 is called a sulfide ion
• The -2 symbol means it
• has gained two electron

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IONIC BONDING

• Cl-1 is called a chloride ion
• The -1 symbol means it
• has gained one electron

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IONIC BONDING

• POLYATOMIC IONS--a group of atoms that act like
  one ion
• NH41--ammonium ion
• CO3-2--carbonate ion
• PO4-3--phosphate ion

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IONIC BONDING

• POLYATOMIC IONS ACT JUST LIKE ANY OTHER NEGATIVE
  ION WHEN BONDING

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IONIC BONDING
SODIUM SULFATE
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IONIC BONDING
3
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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.
• First get them to 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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How to show how they formed

• Its like a jigsaw puzzle.
• I have to tell you what the final formula is.
• You put the pieces together to end up with the
  right formula.
• For example- show how water is formed with
  covalent bonds.

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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 ( I have to tell
  you)
• 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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How to draw them

• Add up all the valence electrons.
• Count up the total number of electrons to make
  all atoms happy.
• Subtract.
• Divide by 2
• Tells you how many bonds - draw them.
• Fill in the rest of the valence electrons to fill
  atoms up.

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Examples
N

• NH3
• N - has 5 valence electrons wants 8
• H - has 1 valence electrons wants 2
• NH3 has 53(1) 8
• NH3 wants 83(2) 14
• (14-8)/2 3 bonds
• 4 atoms with 3 bonds

H
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Examples

• Draw in the bonds
• All 8 electrons are accounted for
• Everything is full

H
N
H
H
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Examples

• HCN C is central atom
• N - has 5 valence electrons wants 8
• C - has 4 valence electrons wants 8
• H - has 1 valence electrons wants 2
• HCN has 541 10
• HCN wants 882 18
• (18-10)/2 4 bonds
• 3 atoms with 4 bonds -will require multiple bonds
  - not to H

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HCN

• Put in single bonds
• Need 2 more bonds
• Must go between C and N

N
H
C
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HCN

• Put in single bonds
• Need 2 more bonds
• Must go between C and N
• Uses 8 electrons - 2 more to add

N
H
C
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HCN

• Put in single bonds
• Need 2 more bonds
• Must go between C and N
• Uses 8 electrons - 2 more to add
• Must go on N to fill octet

N
H
C
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Another way of indicating bonds

• Often use a line to indicate a bond
• Called a structural formula
• Each line is 2 valence electrons

H
H
O
H
H
O

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Structural Examples

• C has 8 electrons because each line is 2
  electrons
• Ditto for N
• Ditto for C here
• Ditto for O

H C N
H
C O
H
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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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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
• d means a partially positive
• d- means a partially negative
  
• 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 ends

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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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MOLECULAR SHAPES

• OF
• COVALENT COMPOUNDS

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VSepR tHEORY
ALENCE
HELL
VSEPR
LECTRON
AIR
EPULSION
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What Vsepr means
Since electrons do not like each other, because
of their negative charges, they orient themselves
as far apart as possible, from each other.
This leads to molecules having specific shapes.
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Things to remember

• Atoms bond to form an Octet (8 outer
  electrons/full outer energy level)
• Bonded electrons take up less space then
  un-bonded/unshared pairs of electrons.

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HERE ARE THE RESULTING MOLECULAR SHAPES
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Linear
EXAMPLE BeF2

• Number of Bonds 2
• Number of Shared Pairs of Electrons 2
• Bond Angle 180

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Trigonal Planar
EXAMPLE GaF3

• Number of Bonds 3
• Number of Shared Pairs of Electrons 3
• Number of Unshared Pairs of Electrons 0
• Bond Angle 120

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Bent 1
EXAMPLE H2O

• Number of Bonds 2
• Number of Shared Pairs of Electrons 2
• Number of Unshared Pairs of Electrons 2
• Bond Angle lt 120

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Bent 2
EXAMPLE O3

• Number of Bonds 2
• Number of Shared Pairs of Electrons 2
• Number of Unshared Pairs of Electrons 1
• Bond Angle gt120

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Tetrahedral
EXAMPLE CH4

• Number of Bonds 4
• Number of Shared Pairs of Electrons 4
• Number of Unshared Pairs of Electrons 0
• Bond Angle 109.5

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Trigonal Pyramidal
EXAMPLE NH3

• Number of Bonds 3
• Number of Shared Pairs of Electrons 4
• Number of Unshared Pairs of Electrons 1
• Bond Angle lt109.5

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Trigonal bIPyramidal
EXAMPLE NbF5

• Number of Bonds 5
• Number of Shared Pairs of Electrons 5
• Number of Unshared Pairs of Electrons 0
• Bond Angle lt120

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OCTAHEDRAL
EXAMPLE SF6

• Number of Bonds 6
• Number of Shared Pairs of Electrons 6
• Number of Unshared Pairs of Electrons 1
• Bond Angle 90

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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

• 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.

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THE END