chapter 5: soap. soap this chapter will introduce the chemistry needed to understand how soap works ...

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