chapter 5: soap. soap this chapter will introduce the chemistry needed to understand how soap works ...
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Chapter 5: Soap
Soap
This chapter will introduce the chemistry needed to understand how soap worksSection 5.1: Types of bondsSection 5.2: Drawing MoleculesSection 5.3: Compounds in 3DSection 5.4: Polarity of MoleculesSection 5.5: Intermolecular ForcesSection 5.6: Intermolecular Forces and
Properties
Section 5.1—Types of Bonds
Why atoms bondAtoms are most stable when they’re outer
shell of electrons is fullAtoms bond to fill this outer shellFor most atoms, this means having 8
electrons in their valence shellCalled the Octet Rule
Common exceptions are Hydrogen and Helium which can only hold 2 electrons
Called the Duet Rule
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
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.
1 2 3 4 5 6 7 8
Remember, Lewis Dot diagrams show their valence electronsWhen 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
LEWIS DOT DIAGRAMS
What is a Bond & Why does it form?
Like glue holding atoms togetherIt’s really forces of attraction between
valence electrons holding atoms togetherThey form because it lowers the potential
energy of the atoms and creates stability
Three Types of Bonding
Ionic Bonding—Metal + 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 electronsThe cation & anion are electrostatically attracted
because of their charges—forming an ionic bond
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.
-
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.
IONIC BONDING
Sodium + Chlorine Sodium Chloride
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
Transfer electrons in ionic bonding
Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge
ClK
Potassium has 1 electron
Chlorine has 7 electrons
Example:Draw the
Lewis Structure for
KCl
+1
Transfer electrons in ionic bonding
Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge
ClK
Potassium has 1 electron
Chlorine has 7 electrons
-1
Example:Draw the
Lewis Structure for
KCl
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
Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
FBa
Barium has 2 electron
Fluorine has 7 electrons
Example:Draw the
Lewis Structure the
ionic compound of
Barium fluoride
The fluorine is full, but the Barium isn’t!
Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
FBa
Barium has 2 electron
Fluorine has 7 electrons
Example:Draw the
Lewis Structure the
ionic compound of
Barium fluoride
F
Add another fluorine atom
Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
FBa
Barium has 2 electron
Fluorine has 7 electrons
Example:Draw the
Lewis Structure the
ionic compound of
Barium fluoride
F
+2 -1
-1
Now all have full valence shells and the charges are balanced, just as when you learned to write in Chpt 2—BaF2!
Covalent BondingWhen two non-metals share electrons2 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
COVALENT BONDING
NONPOLAR COVALENT BONDING
Chlorine + Chlorine Chlorine gas
POLAR COVALENT BONDING
Hydrogen + Flourine Hydrogen Fluoride
+
Metallic Bonding
Metals form a pool of electrons that they share together.
The electrons are free to move throughout the structure—like a sea of electrons
Atoms are bonded as a network
Metallic Bonding
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
Figure 12.4 The three possible types of bonds.
nonpolar
polar
ionic
Sharing electrons equally
Sharing electrons unequally
Transfer of electrons
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
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
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 bondThese compounds are found as solids under normal
conditions
Metals have INTERMEDIATE melting/boiling points. Metallic bonds vary in strength.o All are found as solids under normal conditions except
Hg
Melting/Boiling Points Of Compounds
Covalent Compounds
Polar Covalent molecules have the next highest melting pointsThese compounds are soft solids or liquids under
normal conditions
Non-polar covalent molecules have the lowest melting/boiling pointsThese compounds are found as liquids or gases
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 solventsNonpolar solutes dissolve in nonpolar solvents
Conductivity of Electricity
In order to conduct electricity, charge must be able to move or flow
Metallic bonds have free-moving electrons—they 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
Ionic Compounds ONLY CONDUCT
☺
X
Like Dissolves Like LIke
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.
Na+1 Cl-1
Structures
Metals: sea of electrons
Covalent compounds:separate discrete molecules
Visual Representation of 3 Bonding Types
Section 5.2—Drawing Molecules of Covalent
Compounds
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
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 .
H
H
Example: Carbon Tetrachloride
Example:Draw the
Lewis Structure for
CH4 Remember, “H” can’t go in the middle…put them around the Carbon!
C HH
Carbon has 4 electrons
= 8 electrons
Each hydrogen has 1
H
H
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)
Example:Draw the
Lewis Structure for
CH4
C HHCarbon has 8
Each Hydrogen has 2
All have full valence shells—drawing is correct!
Bonding Pair
Pair of electrons shared by two atoms…they form the “bond”
H
HC HH
Bonding pair
Try These
CH3I PCl3
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
Example:Draw the
Lewis Structure for
CH2O
HC OH
Remember that hydrogen atoms can’t go in the middle!
Example:Draw the
Lewis Structure for
CH2O
HC OH
The two hydrogen atoms are full
But the carbon and oxygen only have 7 each!
Example:Draw the
Lewis Structure for
CH2O
HC OH
But they each have a single, unshared electron.
They could share those with each other!
Example:Draw the
Lewis Structure for
CH2O
HC OH
Now the carbon and oxygen both have a full valence!
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
HC OH
Double Bond
Lone pair
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
…
Example:Draw the
Lewis Structure for
C2H2
HC CH
Remember that hydrogen atoms can’t go in the middle!
…
Example:Draw the
Lewis Structure for
C2H2
HC CH
Each carbon atom only has 7 electrons…not full
Example:Draw the
Lewis Structure for
C2H2
HC CH
But they each have an un-paired electron left!
Example:Draw the
Lewis Structure for
C2H2
HC CH
Now they each have 8 electrons!
Triple Bonds
A Triple Bond occurs when two atoms share 3 pairs of electrons
Triple Bond
HC CH
TRY These:
HCN CO2
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)
Bond Dissociation Energy
Polyatomic Ions
They are a group of atoms bonded together that have an overall charge
Example:Draw the
Lewis Structure for
CO3-2
Drawing Polyatomic Ions
Example:Draw the
Lewis Structure for
CO3-2
C O
When there’s a single atom of one element, put it in the middle
OO
Drawing Polyatomic Ions
Example:Draw the
Lewis Structure for
CO3-2
C O
None of the atoms have full valence shells…they all have 7!
OO
The carbon can double bond with one of the oxygen atoms
Drawing polyatomic Ions
Example:Draw the
Lewis Structure for
CO3-2
C O
Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7
OO
This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!
Drawing Polyatomic Ions
Example:Draw the
Lewis Structure for
CO3-2
C O
Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7
OO
This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!
-2
Covalent bond within…ionic bond between
Polyatomic ions have a covalent bond within themselves…
But can ionic bond with other ions
Covalent bonds within
C OOO
Na
Na
Covalent bond within…ionic bond between
Polyatomic ions have a covalent bond within themselves…
But an ionic bond with other ions
Covalent bonds within
Ionic bond with other ions
C OOO
-2
Na
Na
+1
+1
Section 5.3—Molecules in 3D
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
ValenceShellElectronPairRepulsionTheory
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.
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.
What shapes do molecules form?
Linear
2 bonds, no lone pairs OR any 2 atom molecule
AB2
AB Linear
BeCl2
HCl
What shapes do molecules form?
Trigonal planar
3 bonds, no lone pairs
Indicates a bond coming out at you
Indicates a bond going away from you
AB3
BF3
What shapes do molecules form?
Tetrahedral
4 bonds, no lone pairs
AB4
CH4
What shapes do molecules form?
Trigonal bipyramidal
5 bonds, no lone pairs
AB5
PCl5
What shapes do molecules form?
Octahedral
6 bonds, no lone pairs
AB6
SF6
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!
What shapes do molecules form?
Bent
2 bonds, 1 lone pair
AB2E
AB2E2
SO2
Bent
2 bonds, 2 lone pairs H2O
What shapes do molecules form?
Trigonal pyramidal
3 bonds, 1 lone pair
AB3E
NH3
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.
109.5°
C
105°
O
Example of angle distortion
Rotating Molecular Shapes
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html