Chapter 5: Soap

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Chapter 5 Soap
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Soap

• This chapter will introduce the chemistry needed
  to understand how soap works
• Section 5.1 Types of bonds
• Section 5.2 Drawing Molecules
• Section 5.3 Compounds in 3D
• Section 5.4 Polarity of Molecules
• Section 5.5 Intermolecular Forces
• Section 5.6 Intermolecular Forces and Properties

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Section 5.1Types of Bonds
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Why atoms bond

• Atoms are most stable when theyre outer shell of
  electrons is full
• Atoms bond to fill this outer shell
• For most atoms, this means having 8 electrons in
  their valence shell
• Called the Octet Rule
• Common exceptions are Hydrogen and Helium which
  can only hold 2 electrons
• Called the Duet Rule

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Remember Valence Electrons

• Valence electrons the electrons found in the
  highest energy level
• Short Cut Rule the group next to the letter A
  of the Representative elements represents the
  valence number

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1 How many valence electrons are in an atom?
The main groups of the periodic table each have 1
more valence electron than the group before it.
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2
3
4
5
6
7
8
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Remember, Lewis Dot diagrams show their valence
electrons

• When atoms bond, they have 4 orbitals available
  (1 s and 3 ps). There are 4 places to put
  electrons
• Put one in each spot before doubling up!

Example Draw the Lewis Structure for an oxygen
atom
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LEWIS DOT DIAGRAMS
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What is a Bond Why does it form?

• Like glue holding atoms together
• Its really forces of attraction between valence
  electrons holding atoms together
• They form because it lowers the potential energy
  of the atoms and creates stability

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Three Types of Bonding
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Ionic BondingMetal Non-metal

• Metals have fewer valence electrons and much
  lower ionization energies than non-metals
• Therefore, metals tend to lose their electrons
  and non-metals gain electrons
• Metals become cations (positively charged)
• Non-metals become anions (negatively charged)
• There is a transfer of electrons
• The cation anion are electrostatically
  attracted because of their chargesforming an
  ionic bond

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One way valence shells become full
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Cl
Na
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Sodium has 1 electron in its valence shell
Chlorine has 7 electrons in its valence shell
METALS Some atoms give electrons away to reveal
a full level underneath. NONMETALS Some atoms
gain electrons to fill their current valence
shell.
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One way valence shells become full
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Cl
Na
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The sodium now is a cation (positive charge) and
the chlorine is now an anion (negative
charge). These opposite charges are now
attracted, which is an ionic bond.
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IONIC BONDING

• Sodium Chlorine ? Sodium Chloride

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Transfer electrons in ionic bonding

• Transfer electrons from metal atoms to non-metal
  atoms, keeping track of their new charge

Example Draw the Lewis Structure for KCl
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Transfer electrons in ionic bonding

• Transfer electrons from metal atoms to non-metal
  atoms, keeping track of their new charge

Example Draw the Lewis Structure for KCl
Cl
K
Potassium has 1 electron Chlorine has 7 electrons
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Transfer electrons in ionic bonding

• Transfer electrons from metal atoms to non-metal
  atoms, keeping track of their new charge

Example Draw the Lewis Structure for KCl
-1
1
Cl
K
Potassium has 1 electron Chlorine has 7 electrons
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Add more atoms if needed

• If the transfer from one atom to another doesnt
  result in full outer shells, add more atoms

Example Draw the Lewis Structure the ionic
compound of Barium fluoride
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Add more atoms if needed

• If the transfer from one atom to another doesnt
  result in full outer shells, add more atoms

F
Ba
Example Draw the Lewis Structure the ionic
compound of Barium fluoride
Barium has 2 electron Fluorine has 7 electrons
The fluorine is full, but the Barium isnt!
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Add more atoms if needed

• If the transfer from one atom to another doesnt
  result in full outer shells, add more atoms

F
Ba
Example Draw the Lewis Structure the ionic
compound of Barium fluoride
F
Barium has 2 electron Fluorine has 7 electrons
Add another fluorine atom
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Add more atoms if needed

• If the transfer from one atom to another doesnt
  result in full outer shells, add more atoms

-1
2
F
Ba
Example Draw the Lewis Structure the ionic
compound of Barium fluoride
-1
F
Barium has 2 electron Fluorine has 7 electrons
Now all have full valence shells and the charges
are balanced, just as when you learned to write
in Chpt 2BaF2!
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Covalent Bonding

• When two non-metals share electrons
• 2 Identical Non-metals that share electrons
  evenly form non-polar covalent bonds
• 2 Different Non-metals that share electrons
  un-evenly form polar covalent bonds

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COVALENT BONDING

• NONPOLAR COVALENT BONDING
• Chlorine Chlorine ? Chlorine gas
• POLAR COVALENT BONDING
• Hydrogen Flourine ? Hydrogen Fluoride
• ?

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Metallic Bonding

• Metals form a pool of electrons that they share
  together.
• The electrons are free to move throughout the
  structurelike a sea of electrons
• Atoms are bonded as a network

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Metallic Bonding
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Bonding in Never Purely Ionic or Covalent Use EN
to indicate primary type

• 0 5 50
  100
• 0-----.3 --------------------1.7------------------
  -3.3
• np polar
  ionic
• covalent covalent
• If the electronegativity difference is
• Greater than 1.7 IONIC
• Between 0-.29 NONPOLAR COVALENT
• Between .3 and 1.69 POLAR COVALENT

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Figure 12.4 The three possible types of bonds.
nonpolar
Sharing electrons equally
polar
Sharing electrons unequally
ionic
Transfer of electrons
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Examples

• CH4
• F2
• H2O
• NaF

2.5 2.1 .4 polar covalent
4.0 -4.0 0 nonpolar covalent
3.5 2.1 1.4 polar covalent
4.0 .9 3.1 ionic
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Bond type affects properties

• There are always exceptions to these
  generalizations (especially for very small or
  very big molecules), but overall the pattern is
  correct

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Melting/Boiling Points Of Compounds

• Ionic Compounds tend to have very high
  melting/boiling points as its hard to pull apart
  those electrostatic attractions of the ionic bond
• These compounds are found as solids under normal
  conditions
• Metals have INTERMEDIATE melting/boiling points.
  Metallic bonds vary in strength.
• All are found as solids under normal conditions
  except Hg

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Melting/Boiling Points Of Compounds

• Covalent Compounds
• Polar Covalent molecules have the next highest
  melting points
• These compounds are soft solids or liquids under
  normal conditions
• Non-polar covalent molecules have the lowest
  melting/boiling points
• These compounds are found as liquids or gases

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Solubility in Water

• Ionic polar covalent compounds tend to be
  soluble in water
• Non-polar metallic compounds tend to be
  insoluble in water
• Use this Like Dissolves Like rule of thumb.
• Polar solutes dissolve in polar solvents
• Nonpolar solutes dissolve in nonpolar solvents

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Conductivity of Electricity

• In order to conduct electricity, charge must be
  able to move or flow
• Metallic bonds have free-moving electronsthey
  can conduct electricity in solid and liquid state
• Ionic bonds have free-floating ions when
  dissolved in water or in liquid form that allow
  them conduct electricity
• Covalent bonds never have charges free to move
  and therefore cannot conduct electricity in any
  situation

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Ionic Compounds ONLY CONDUCT
X
?
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Like Dissolves
Like LIke
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Structures Ionic Compounds

• Ionic compounds are made of positive and negative
  ions.
• They pack together so that the like-charge
  repulsions are minimized while the
  opposite-charge attractions are enhanced.

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Structures

• Metals sea of electrons
• Covalent compoundsseparate discrete molecules

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Visual Representation of 3 Bonding Types
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Section 5.2Drawing Molecules of Covalent
Compounds
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Tips for arranging atoms

• In general, write out the atoms in the same order
  as they appear in the chemical formula
• Hydrogen Halogens (F, Cl, Br, I) can only bond
  with one other atom
• Always put them around the outside
• The least electronegative atom is usually in the
  middle Carbon always goes in the middle

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Steps to Drawing Lewis Structures

• 1. Decide how many valence electrons are around
  each atom
• 2. Arrange the atoms in a skeletal structure and
  connect them with a bonding pair of electrons.
• 3. Place remaining electrons around atoms so
  they each acquire 8 electrons. Exception is H .

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Example Carbon Tetrachloride
Carbon has 4 electrons
8 electrons Each hydrogen has 1
Example Draw the Lewis Structure for CH4

Remember, H cant go in the middle put them
around the Carbon!
H
H
H
C
H
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Count electrons around each atom

• Any electron that is being shared (between two
  atoms) gets to be counted by both atoms!
• All atoms are full with 8 valence electrons
  (except Hcan only hold 2)

H
Example Draw the Lewis Structure for CH4
Carbon has 8
H
H
C
Each Hydrogen has 2
H
All have full valence shellsdrawing is correct!
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Bonding Pair

• Pair of electrons shared by two atomsthey form
  the bond

Bonding pair
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Try These

• CH3I PCl3

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What if theyre not all full after that? Multiple
Covalent Bonds

• Are needed when there is not enough electrons to
  complete an octet
• To satisfy move lone pair in between atoms to
  satisfy the duet/octet rule

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H
Example Draw the Lewis Structure for CH2O
O
H
C
Remember that hydrogen atoms cant go in the
middle!
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The two hydrogen atoms are full
But the carbon and oxygen only have 7 each!
H
Example Draw the Lewis Structure for CH2O
O
H
C
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But they each have a single, unshared
electron. They could share those with each other!
H
Example Draw the Lewis Structure for CH2O
O
H
C
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Now the carbon and oxygen both have a full
valence!
H
Example Draw the Lewis Structure for CH2O
O
H
C
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Double Bonds Lone Pairs

• Double bonds are when 2 pairs of electrons are
  shared between the same two atoms
• Lone pairs are a pair of electrons not
  sharedonly one atom counts them

Lone pair
Double Bond
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And when a double bond isnt enough

• Sometimes forming a double bond still isnt
  enough to have all the valence shells full

Example Draw the Lewis Structure for C2H2
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Example Draw the Lewis Structure for C2H2
H
C
H
C
Remember that hydrogen atoms cant go in the
middle!
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Each carbon atom only has 7 electronsnot full
Example Draw the Lewis Structure for C2H2
H
C
H
C
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But they each have an un-paired electron left!
Example Draw the Lewis Structure for C2H2
H
C
H
C
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Now they each have 8 electrons!
Example Draw the Lewis Structure for C2H2
H
C
H
C
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Triple Bonds

• A Triple Bond occurs when two atoms share 3 pairs
  of electrons

Triple Bond
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TRY These

• HCN CO2

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Properties of multiple bonds

• Single Bond
• Double Bond
• Triple Bond

Shorter (atoms closer together Stronger
bonds (Higher Bond Energy takes more energy to
break)
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Bond Dissociation Energy
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Polyatomic Ions

• They are a group of atoms bonded together that
  have an overall charge

Example Draw the Lewis Structure for CO3-2
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Drawing Polyatomic Ions
O
Example Draw the Lewis Structure for CO3-2
O
C
O
When theres a single atom of one element, put it
in the middle
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Drawing Polyatomic Ions
None of the atoms have full valence shellsthey
all have 7!
The carbon can double bond with one of the oxygen
atoms
O
Example Draw the Lewis Structure for CO3-2
O
C
O
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Drawing polyatomic Ions
Now the Carbon and the one oxygen have 8but the
other two oxygen atoms still only have 7
O
Example Draw the Lewis Structure for CO3-2
O
C
O
This is a polyatomic ion with a charge of
-2that means we get to add 2 electrons!
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Drawing Polyatomic Ions
Now the Carbon and the one oxygen have 8but the
other two oxygen atoms still only have 7
-2
O
Example Draw the Lewis Structure for CO3-2
O
C
O
This is a polyatomic ion with a charge of
-2that means we get to add 2 electrons!
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Covalent bond withinionic bond between

• Polyatomic ions have a covalent bond within
  themselves
• But can ionic bond with other ions

Na
O
O
C
O
Covalent bonds within
Na
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Covalent bond withinionic bond between

• Polyatomic ions have a covalent bond within
  themselves
• But an ionic bond with other ions

1
-2
Na
O
O
C
Ionic bond with other ions
O
Covalent bonds within
1
Na
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Section 5.3Molecules in 3D
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Bonds repel each other
Bonds are pairs electrons. Electrons are
negatively charged
Negative charges repel other negative charges
SO bonds repel each other
Molecules arrange themselves in 3-D so that the
bonds are as far apart as possible
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Valence Shell Electron Pair Repulsion Theory
(VSEPR Theory)
Valence Shell Electron Pair Repulsion Theory
Outer shell of electrons involved in bonding
Bonds are made of electron pairs
Those electron pairs repel each other
Attempts to explain behavior
This theory attempts to explain the 3-D shape of
molecules.
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Coding for a Shape

• A code can help you connect the molecule to its
  shape.
• A stands for the central atom
• B stands for the number of bonding atoms off
  the central atom.
• E stands for the number of lone pair coming off
  the CENTRAL atom.

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What shapes do molecules form?
AB2
BeCl2
Linear
AB
HCl
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What shapes do molecules form?
AB3
BF3
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What shapes do molecules form?
CH4
AB4
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What shapes do molecules form?
PCl5
AB5
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What shapes do molecules form?
SF6
AB6
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Lone Pairs

• Lone pairs are electrons, toothey must be taken
  into account when determining molecule shape
  since they repel the other bonds as well.
• But only take into account lone pairs around the
  CENTRAL atom, not the outside atoms!

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What shapes do molecules form?
SO2
AB2E
H2O
AB2E2
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What shapes do molecules form?
NH3
AB3E
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Lone Pairs take up more space

• Lone pairs arent controlled by a nucleus
  (positive charge) on both sides, but only on one
  side.
• This means they spread out more than a bonding
  pair.
• They distort the angle of the molecules bonds
  away from the lone pair.

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Example of angle distortion
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Rotating Molecular Shapes

• http//intro.chem.okstate.edu/1314F00/Lecture/Chap
  ter10/VSEPR.html