1. Unit 5: Bonding

2. 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

3. Section 1—Types of Bonds

4. Why do atoms bond?
Atoms are most stable when they’re 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

5. 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

6. LEWIS DOT DIAGRAMS

7. Remember, Lewis Dot diagrams show their valence electrons
When atoms bond, they have 4 orbitals available (1 “s” and 3 “p’s”). There are 4 places to put electrons
Put one in each spot before doubling up!
Example:
Draw the Lewis Structure for an oxygen atom

8. What is a Bond & Why does it form?
Like glue holding atoms together
It’s really forces of attraction between valence electrons holding atoms together
They form because it lowers the potential energy of the atoms and creates stability

9. Three Types of Bonding

10. Ionic Bonding—Metal + Non-metal
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 charges—forming an ionic bond

12. One way valence shells become full- - - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - -
Sodium has 1 electron in it’s valence shell
Chlorine has 7 electrons in it’s 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.

13. One way valence shells become full- - - + - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - -
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.

14. IONIC BONDING
Sodium + Chlorine  Sodium Chloride

15. 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

16. 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 K Cl
Potassium has 1 electron
Chlorine has 7 electrons

17. Transfer electrons in ionic bonding
Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge
Potassium has 1 electron
Chlorine has 7 electrons
Example:
Draw the Lewis Structure for KCl +1 -1 K Cl

18. Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
Example:
Draw the Lewis Structure the ionic compound of Barium fluoride

19. Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
Example:
Draw the Lewis Structure the ionic compound of Barium fluoride Ba F
Barium has 2 electron
Fluorine has 7 electrons
The fluorine is full, but the Barium isn’t!

20. Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
Example:
Draw the Lewis Structure the ionic compound of Barium fluoride Ba F F
Barium has 2 electron
Fluorine has 7 electrons
Add another fluorine atom

21. Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms -1 +2
Example:
Draw the Lewis Structure the ionic compound of Barium fluoride Ba F -1 F
Barium has 2 electron
Fluorine has 7 electrons
Now all have full valence shells and the charges are balanced—BaF2!

22. 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

23. COVALENT BONDING
NONPOLAR COVALENT BONDING Chlorine + Chlorine  Chlorine gas POLAR COVALENT BONDING Hydrogen + Fluorine  Hydrogen Fluoride + 

25. Metallic Bonding: composed only of metals
Metal atom nuclei are bonded as a network as a pool of electrons that they share together.
The electrons are free to move(mobile) throughout the structure—like a sea of electrons

26. Metallic Bonding

27. Bonding in Never Purely Ionic or Covalent: Use EN to indicate primary type
0% 5% % %
np polar ionic
covalent covalent
If the electronegativity difference is:
Between NONPOLAR COVALENT
Between .41 and POLAR COVALENT
Greater than IONIC

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

29. CH4 F2 H2O NaF NM/NM M/NM Examples 2.5 – 2.1 = .4 non-polar covalent
= 0
nonpolar covalent
3.5 – 2.1 = 1.4
polar covalent
4.0 – .9 = 3.1
ionic

30. 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

31. Melting/Boiling Points Of Compounds
Ionic Compounds tend to have VERY HIGH melting/boiling points as it’s hard to pull apart those electrostatic attractions of the ionic bond. Strongest forces of attraction.
These compounds are found as solids under normal conditions
Metals have HIGH/INTERMEDIATE melting/boiling points. Metallic bonds vary in strength. In general, strong forces of attraction.
All are found as solids under normal conditions except Hg

32. Melting/Boiling Points Of Compounds
Covalent Compounds
Polar Covalent molecules have the next highest melting points. Weak forces of attraction.
These compounds are soft solids or liquids under normal conditions
Non-polar covalent molecules have the lowest melting/boiling points. Weakest forces of attraction.
These compounds are found as liquids or gases

33. Solubility in Water (polar molecule)
Ionic & polar covalent compounds tend to be soluble in water
Non-polar compounds & metals tend to be insoluble in water
Use this: Like Dissolves Like rule of thumb.
Polar dissolves Polar
Nonpolar dissolves Nonpolar

34. Like Dissolves Like

35. Conductivity of Electricity
In order to conduct electricity, charge must be able to move or flow
Metals have free-moving electrons—they can conduct electricity in BOTH solid & liquid state
Ionic compounds ONLY have free-floating ions when dissolved in water, aqueous form or when melted (liquid state)
Covalent compounds never have charges free to move and therefore cannot conduct electricity
Exception:
POLAR COVALENT Molecules of acids & bases
can conduct electricity when dissolved in water

36. Ionic Compounds ONLY CONDUCTX ☺

37. Structures: Ionic Compounds(Crystals)
Ionic compounds form a lattice framework of positive and negative ions
They pack together so that the like-charge repulsions are minimized while the opposite-charge attractions are enhanced.
Na+1
Cl-1

38. Metals: sea of electrons
Structures
Metals: sea of electrons
Covalent compounds: neutral, separate discrete molecules

39. Visual Representation of 3 Bonding Types

40. Section 2—Drawing Molecules of Covalent Compounds

41. 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

42. 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 .

43. Example: Carbon Tetrachloride
Carbon has 4 electrons
= 8 electrons
Each hydrogen has 1
Example:
Draw the Lewis Structure for CH4
Remember, “H” can’t go in the middle…
put them around the Carbon! H H C H H

44. 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 H—can only hold 2) H
Example:
Draw the Lewis Structure for CH4
Carbon has 8 H C H
Each Hydrogen has 2 H
All have full valence shells—drawing is correct!

45. Bonding Pair
Pair of electrons shared by two atoms…they form the “bond” H C
Bonding pair

46. Try These
CH3I PCl3

47. What if they’re 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

48. Examples with Double Bonds
Draw the Lewis Structure for CH2O H C O
Remember that hydrogen atoms can’t go in the middle!

49. Draw the Lewis Structure for CH2O
The two hydrogen atoms are full
But the carbon and oxygen only have 7 each! H
Example:
Draw the Lewis Structure for CH2O H C O

50. Draw the Lewis Structure for CH2O
But they each have a single, unshared electron.
They could share those with each other! H
Example:
Draw the Lewis Structure for CH2O H C O

51. Draw the Lewis Structure for CH2O
Now the carbon and oxygen both have a full valence! H
Example:
Draw the Lewis Structure for CH2O H C O

52. 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 shared—only one atom “counts” them H C O
Lone pair
Double Bond

53. And when a double bond isn’t enough…
Sometimes forming a double bond still isn’t enough to have all the valence shells full
Example:
Draw the Lewis Structure for C2H2

54. Draw the Lewis Structure for C2H2…
Example:
Draw the Lewis Structure for C2H2 H C C H
Remember that hydrogen atoms can’t go in the middle!

55. Draw the Lewis Structure for C2H2…
Each carbon atom only has 7 electrons…not full
Example:
Draw the Lewis Structure for C2H2 H C C H

56. Draw the Lewis Structure for C2H2
But they each have an un-paired electron left!
Example:
Draw the Lewis Structure for C2H2 H C C H

57. Draw the Lewis Structure for C2H2
Now they each have 8 electrons!
Example:
Draw the Lewis Structure for C2H2 H C C H

58. Triple Bonds
A Triple Bond occurs when two atoms share 3 pairs of electrons H C
Triple Bond

59. TRY These:
HCN CO2

60. Properties of multiple bonds
Single Bond
Double Bond
Triple Bond
Longer & Weaker bonds
(Lower Bond Energy: takes less
energy to break)
Shorter (atoms closer together &
Stronger bonds
(Higher Bond Energy: takes more
energy to break)

61. Bond Dissociation Energy

62. Draw the Lewis Structure for CO3-2
Polyatomic Ions
They are a group of atoms bonded together that have an overall charge
Example:
Draw the Lewis Structure for CO3-2

63. Drawing Polyatomic Ions
Example:
Draw the Lewis Structure for CO3-2 O C O O
When there’s a single atom of one element, put it in the middle

64. Drawing Polyatomic Ions
None of the atoms have full valence shells…they all have 7!
The carbon can double bond with one of the oxygen atoms
Example:
Draw the Lewis Structure for CO3-2 O C O O

65. Drawing polyatomic Ions
Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7
Example:
Draw the Lewis Structure for CO3-2 O C O O
This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!

66. Drawing Polyatomic Ions
Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7 -2
Example:
Draw the Lewis Structure for CO3-2 O C O O
This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!

67. Covalent bond within…ionic bond between
Polyatomic ions have a covalent bond within themselves…
But can ionic bond with other ions Na O C O O
Covalent bonds within Na

68. Covalent bond within…ionic bond between
Polyatomic ions have a covalent bond within themselves…
But an ionic bond with other ions +1 -2 Na O C O
Ionic bond with other ions O
Covalent bonds within +1 Na

69. Section 3—Molecules in 3D

70. 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

71. Valence Shell Electron Pair Repulsion Theory (VSEPR 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.

72. 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.

73. What shapes do molecules form?
Linear
2 bonds, no lone pairs OR any 2 atom molecule
AB2
BeCl2
NO CODE: a 2 atom molecule has no central atom
Linear
HCl
180°

74. What shapes do molecules form?
Trigonal planar
3 bonds, no lone pairs
AB3
BF3
Indicates a bond going away from you
120°
Indicates a bond coming out at you

75. What shapes do molecules form?
Tetrahedral
4 bonds, no lone pairs
CH4
AB4
109.5°

76. What shapes do molecules form?
Trigonal bipyramidal
5 bonds, no lone pairs
PCl5
AB5
90° and 120°

77. What shapes do molecules form?
Octahedral
6 bonds, no lone pairs
SF6
AB6
90°

78. How is shape affected by Lone Pairs?
Lone pairs are electrons, too…they 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!

79. What shapes do molecules form?
Bent
2 bonds, 1 lone pair
SO2
AB2E
Bent
2 bonds, 2 lone pairs
H2O
AB2E2

80. What shapes do molecules form?
NH3
Trigonal pyramidal
3 bonds, 1 lone pair
AB3E

81. Lone Pairs take up more space
Lone pairs aren’t “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 molecule’s bonds away from the lone pair.

82. Example of angle distortion
105° O
109.5° C

84. Rotating Molecular Shapes

85. Section 4—Polarity of Molecules

86. Showing Partial Charges of POLAR BONDS
Determine which element in the bond is more electronegative. Use the symbol “-” to express the more electronegative atom and the symbol “+” for the least electronegative atom
Use of an arrow pointing towards the partial negative atom with a “plus” tail at the partial positive atom. The arrow represents DIPOLES which are: forces created due to a separation of opposite charge!( a difference in charged ends) C O H + - C O H

87. Let’s Practice: If the bond is polar, draw the polarity arrow
C – H
O—Cl
F—F
P—N

88. If the bond is polar, draw the polarity arrow
Let’s Practice
C – H
O—Cl
F—F
P—N
2.5 – 2.1 = non-polar*exception
3.5 – 3.0 = polar
4.0 – 4.0 = non-polar
3.0 – 2.1 = polar
Example:
If the bond is polar, draw the polarity arrow

89. Predicting Bond Polarity Two atoms sharing equally: Draw N2
Each nitrogen atom has an electronegativity of They pull evenly on the shared electrons. The electrons are not closer to one or the other of the atoms. This is a non-polar covalent bond.
Most compounds that contain ONLY non-polar covalent bonds are NONPOLAR MOLECULES!

90. Atoms sharing unequally: Draw H2S
Electronegativities: H = sulfur = 2.5
The sulfur pulls on the electrons slightly more, pulling them slightly towards the sulfur.
This is a polar covalent bond
Asymmetrical (lop-sided) compounds that
contain polar covalent bonds are
POLAR MOLECULES

91. Sharing unevenly: Draw CH2O
Electronegativities: H = C = 2.5
O = 3.5
The carbon-hydrogen difference isn’t great enough to create partial charges : It’s actually a NON POLAR bond: **Exception to the rule
But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond.
Asymmetrical (lop-sided) compounds that contain
polar covalent bonds are POLAR MOLECULES

92. Sharing unevenly: Draw BCl3
Electronegativities:
B = Cl = 3.0
The chlorine pulls on the electrons slightly more than boron This is a polar covalent bond
Symmetrical compounds that contain polar covalent bonds are NONPOLAR MOLECULES

93. How do Dipoles Cancel?
Dipoles must move in equal but opposite directions in order for the forces to cancel The molecule is classified as NONPOLAR.

94. Polar Bonds versus Polar Molecules
Not every molecule with a polar bond is polar itself
If the polar bonds form Dipoles that cancel out then the molecule is overall non-polar.
The dipoles cancel out.
No net dipole
The dipoles do
not cancel out.
Net dipole
This one is hard to tell!

95. The Importance of VSEPR in Predicting Polarity.
Shape is important. All molecules must be drawn in the correct shape to see the proper canceling of dipoles to determine if its polar or nonpolar.
Water drawn this way shows all the dipoles canceling out. O H
But water drawn in the correct VSEPR structure, bent, shows the dipoles don’t cancel out!
Net dipole
H O H

96. Practice: Draw the molecule NH3
Example:
Is NH3 a polar molecule?

97. N H Example: Is NH3 a polar molecule? Electronegativities: N = 3.0
Difference = Polar bonds
VSEPR shape = Trigonal pyramidal
Yes, NH3 is polar
Net dipole

98. Draw the molecule for dihydrogen monoxide, H2O.
Is it polar or non-polar? What shape?

99. O H Net dipole Yes, H2O is polar
Is water polar or non-polar? What shape?
Electronegativities:
H = 2.1
O = 3.0
Difference = .9 Polar bonds
VSEPR shape = bent

100. Draw the molecule CO2
Example:
Is CO2 a polar molecule?

101. Draw the molecule of carbon dioxide, CO2
Example:
Is CO2 a polar molecule?
Electronegativities:
C = 2.5
O = 3.5
Difference = Polar bonds
VSEPR shape = linear
No CO2 is nonpolar
Dipole cancels

102. Draw the molecule for carbon tetrachloride, CCl4.
Electronegativities:
C = 2.5
H = 2.1
Difference = .4 Nonpolar bonds
VSEPR shape = tetrahedral H H C H H
No Net dipole
Yes, CH4 is nonpolar

103. Let’s make this Simple:
Nonpolar bonds= Nonpolar molecule.
Polar bonds with a lone pair on the central atom = most likely a polar molecule
Polar bonds & no lone pair on central atom & all terminal atoms are the same= nonpolar molecule.
If terminal atoms are different, its polar.

105. Section 5—Intermolecular Forces

106. Intramolecular Forces- versus Inter-molecular Forces
So far this chapter has been discussing “Intramolecular Forces”
Intramolecular forces = forces within the molecule
AKA: chemical bonds

107. Intra-Force
Intra-Force

108. Breaking Intramolecular forces
Breaking of intramolecular forces (within the molecule) is a chemical change
Example: H2 + O2  2 H2O
Bonds are broken within the molecules and new bonds are formed to form new molecules
Requires a larger amount of energy to break than an intermolecular force

109. Inter-molecular Forces
Intermolecular forces = forces between separate molecules

110. Intra-Force
Inter-Force

111. Breaking Intermolecular forces
Breaking of intermolecular forces (between separate molecules) is a physical change
Breaking glass & Boiling water are examples
Example: H2O(l)  2 H2O(g)
Does not require as much energy to break compared to an intramolecular force

112. London Dispersion Forces (LDF): are the primary force between nonpolar molecules but are found in all molecules!
Electrons move
around the nuclei &
They can momentarily
“gang up” on one side
This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another. +
Positively charged nucleus -
Negatively charged electron + - + -
Electrons are fairly evenly dispersed. + -
As electrons move, they “gang up” on one side.

113. London Dispersion Forces (LDF)
Once the electrons have “ganged up”, a temporary dipole is created & the molecule is now temporarily polar.
The positive end of one temporarily polar molecule can be attracted to the negative end of another molecule. + - + -
LDF
nonpolar
nonpolar

114. London Dispersion Forces (LDF)

115. Strength of London Dispersion Forces (LDF)
Since, the electrons will move again, returning the molecule back to non-polar
The polarity was temporary, therefore the molecule cannot always form LDF.
London Dispersion Forces are the weakest IMF because molecules can’t form it all the time, only temporarily

116. Strength of London Dispersion Forces (LDF)
All molecules have electrons…all molecules can have London Dispersion Forces
Larger molecules (greater molar mass) have stronger London Dispersion Forces than smaller molecules.
Larger molecules have more electrons.
The more electrons that gang-up, the larger the partial negative charge.

117. London Forces explain why Chlorine is a gas, Bromine is a liquid and Iodine is a solid!
Chlorine Gas = 34 e-
Bromine Liquid = 70 e-
GREATER # ELECTRONS,
Greater Molar Mass, STRONGER FORCES
Iodine Solid = 106 e-

118. Dipole- Dipole Forces: primary force between polar molecules
They have permanent permanent dipoles.
The positive area of one polar molecule can be attracted to the negative area of another molecule.
The partial positive & negative poles are shown as + and -

119. Strength of Dipole Forces
In general, Dipole forces are stronger than London Dispersion Forces
Polar molecules always have a partial separation of charge.
They always have the ability to form attractions with opposite charges + - + -
Dipole force
polar
polar

120. Dipole-Dipole Forces

121. Hydrogen Bonding
A special dipole force between polar molecules that contain a hydrogen atom connecting with F, O, or N of another molecule.
(H is FON)
A very strong dipole forms since F, O, and N are all very small, highly electronegative atoms.

122. Hydrogen Bond N H N H
Hydrogen bond

123. Hydrogen Bonds
The ladder rungs in a DNA molecule are hydrogen bonds between the base pairs, (AT and GC).

124. Special Case: Carbon Allotropes: Diamond vs. Graphite
Diamond: Hard
Tetrahedral-Special : Network Covalent Bonds
Graphite: soft
Strong Sheets of carbon rings but weak forces holding the sheets together

125. NETWORK COVALENT BONDS
special covalent compounds
compounds that contain only carbon (diamond, graphite) or silicon compounds (silicon dioxide- quartz)
super strong bonding
super high melting points

126. Rank the forces of attraction in order of weakest to strongest
Rank the Intramolecular Forces: Ionic, Covalent, and Metallic
Covalent< Metallic < Ionic
Rank the Intermolecular Forces: Dipole, London Dispersion, Hydrogen bonding
London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding
Rank ALL the Forces:
London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding <Covalent< Metallic< Ionic , NETWORK COVALENT

127. Tutorial: must be in Mozilla
Wisconsin online :intermolecular forces

128. Section.6—Intermolecular Forces & Properties

129. IMF’s and Properties
The number and strength of the intermolecular forces affects the properties of the substance.
Energy is needed to break IMF’s
Energy is released when new IMF’s are formed

130. IMF’s and Changes in State
From solid to liquid, some IMF’s are broken.
From liquid to gas, the remainder of the IMF’s are broken
Breaking IMF’s requires energy =
ENDOTHERMIC
PROCESS
The stronger the IMF’s are, the higher the
melting and boiling point

131. The Uniqueness of Water
Water is a very small molecule
In general small molecules have low melting and boiling points
Based on it’s size, water should be a gas under normal conditions
However, water is polar and can form dipole interactions and hydrogen bonding,
It’s boiling point is much higher as a result!

132. Boiling Point of Polar Molecules

133. What is More Viscous? Molasses or Water?

134. Viscosity Viscosity is the resistance to flow
Molasses is much more viscous than
water
Larger molecules and molecules with strong IMF’s become inter-twined and “stick” together more
More sticking together = higher viscosity
An increase in temperature will help break the IMF’s and make a substance less viscous

135. Solubility Solute: the substance that is dissolved Solvent: the
is doing the
dissolving

136. Solubility: the ability to dissolve
Solvent, water (polar) + -
- +
Solute, sugar (polar)
Water particles break IMF’s between other water molecules allowing them to spread out to form new IMF’s with the sugar molecules.
- +
- +
- +

137. Solubility
Solvent, water (polar) + -
- +
Solute, sugar (polar)
As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving”
- +

139. Solubility
If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occur
An exception to this is if more energy is added somehow (such as heating)

140. Like Dissolves Like Polar solvents dissolve polar solutes
Nonpolar solvents dissolve nonpolar solutes
Polar solvents can also dissolve ionic compounds because of the charged ends of both

141. Therefore, oil and water don’t mix!
Oil & Water
Water is polar and can hydrogen bond, Oil is non-polar.
Water can only form London Dispersion with the oil. That doesn’t release much energy
Much more energy is required to break apart the water than is released when water and oil combine.
Water has London Dispersion, Dipole forces and hydrogen bonding. That takes a lot of energy to break
Therefore, oil and water don’t mix!

142. Surface Tension
Surface tension is the resistance of a liquid to spread out.
Due to higher IMF’s in the liquid, the more the molecules “stick” together, the less they want to spread out.
Higher IMF’s =higher the surface tension.

143. Surface Tension of Water
metal paper clip on water
water forms “beads”