chapter 5: soap. introductory activity fill a test tube with an inch of water add a squirt of...

Chapter 5: Soap

Introductory Activity

Fill a test tube with an inch of waterAdd a squirt of cooking oil to the test tube.

ObserveStopper, shake & observeAdd a few drops of soap. ObserveStopper, shake & observeWith another test tube, add water & soap only.

Observe.Compare the two test tubes.Make particle visualizations describing each test

tube.

Introductory Activity

What ideas do you have about how soap works?

What kinds of things do advertising and marketing tell you?

What do the soap companies want you to know about how soap works?

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

Soap

Inter-molecular forces


Works based on

Molecular Geometry

Molecular Geometry

Bonding types &

Structures

Bonding types &

Structures

Determined by

Determined by

Section 5.1—Types of Bonds

Why atoms bond

Atoms are most stable when they’re outer shell of electrons is full

Atoms bonds 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.

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

Some atoms give electrons away to reveal a full level underneath.

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—Metal + Non-metal

Metals have fewer valence electrons and much lower ionization energies (energy needed to remove an electron) 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)The cation & anion are attracted because of

their charges—forming an ionic bond

Bonding between non-metals

When two non-metals bond, neither one loses or gains electrons much more easily than the other one.

Therefore, they share electronsNon-metals that share electrons evenly

form non-polar covalent bondsNon-metals that share electrons un-evenly

form polar covalent bonds

Metals 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 aren’t bonded to specific other atoms, but rather to the network as a whole

Bond type affects properties

The type of bonding affects the properties of the substance.

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

Melting/Boiling Points

Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractionsThey’re found as solids under normal conditions

Polar covalent bonds have the next highest melting/boiling pointsMost are solids or liquids under normal conditions

Non-polar covalent bonds have lower melting/boiling pointsMost 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

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

Section 5.2—Drawing Molecules

Drawing Molecules on Paper

Lewis Structures (or Dot Structures) are one way we draw molecules on paper

Since paper is 2-D and molecules aren’t, it’s not a perfect way to represent how molecules bond…but it’s a good way to begin to visualize molecules

Drawing Ionic Compounds

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

2: Placing electrons around an atom

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

3: 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

4: 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

4: 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

A note about Ionic Dot Structures

The atoms are not sharing the electrons—make sure you clearly draw the atoms separate!

Drawing Covalent Compounds

Tips for arranging atoms

Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom—they can’t go in the middle of a molecules

Always put them around the outside

In general, write out the atoms in the same order as they appear in the chemical formula

Repeat first two steps from before

1. Use the periodic table to decide how many electrons are around each atom

2. Write the electrons around each atom

Example:Draw the

Lewis Structure for

CH4

H

H

Repeat first two steps from before

1. Use the periodic table to decide how many electrons are around each atom

2. Write the electrons around each atom

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

Each hydrogen has 1

H

H

3: 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

What if they’re not all full after that?

Sometimes, the first 3 steps don’t leave you with full valence shells for all atoms

Example:Draw the

Lewis Structure for

CH2O

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

Properties of multiple bonds

Single Bond

Double Bond

Triple Bond

Shorter bonds (atoms closer together)

Stronger bonds (takes more energy to break)

Polyatomic Ions

Polyatomic Ions

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

Example: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

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 an ionic bond with other ions

Covalent bonds within

Ionic bond with other ions

C OOO

-2

Na

Na

+1

+1

Isomers

More than one possibility

Often, there’s more than one way to correctly draw a Dot Structure

HC CH CHH

HC CH CH

H

Chemical Formula: C3H4


Contains 2 sets of double bonds between carbons

Contains 1 triple bond and 1 single bond between carbons

Both structures have full valence shells!

Both are “correct”

The chemical formula alone does not give you enough information to differentiate between the two structures

HC CH CHH

HC CH CH

H


You’ll learn in Chapter 11 how to differentiate between these two structures

with chemical names

Isomers

Isomers: Structures with the same chemical formula but different chemical structure

Atoms must be bonded differently (multiple versus single bonds) or in a different order) but have the same overall chemical formula to be isomeric structures

Section 5.3—Molecules in 3D

Bonds repel each other

Bonds are electrons. Electrons are negatively charged

Negative charges repel other negative charges

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 (that bonds repel each other because they’re like charges) attempts to explain why molecules form the shapes they form

What shapes do molecules form?

Linear

2 bonds, no lone pairs

Trigonal planar3 bonds, no lone pairs

Indicates a bond coming out at you

Indicates a bond going away from you


Tetrahedron


Trigonal bipyramidal



Octahedron


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!


Bent

2 bonds, 1 lone pair

Trigonal pyramidal

3 bonds, 1 lone pair


Bent

2 bonds, 2 lone pairs

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

Ionic Compound structures

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

Section 5.4—Polarity of Molecules

Electronegativity

The pull an atom has for the electrons it shares with another atom in a bond.

Electronegativity is a periodic trendAs atomic radius increases and number of

electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons

Periodic Table with Electronegativies

increases

decreases

Polar Bond

A polar covalent bond is when there is a partial separation of charge

One atom pulls the electrons closer to itself and has a partial negative charge.

The atom that has the electrons farther away has a partial positive charge

Two atoms sharing equally

N N

Each nitrogen atom has an electronegativity of 3.0

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

Atoms sharing almost equally

Electronegativities: H = 2.1 C = 2.5

The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon

Put the difference isn’t enough to create a polar bond

This is a non-polar covalent bond

C HH

H

H

Sharing unevenly

Electronegativities: H = 2.1 C = 2.5 O = 3.5

The carbon-hydrogen difference isn’t great enough to create partial charges

But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond

This is a polar covalent bond

C OH

H

Showing Partial Charges

There are two ways to show the partial separation of chargesUse of “” for “partial” Use of an arrow pointing towards the partial

negative atom with a “plus” tail at the partial positive atom

C OH

H

+ -C OH

H

Ionic Bonds

Ionic bonds occur when the electronegativies of two atoms are so different that they can’t even share unevenly…one atom just takes them from the other

How to determine bond type

Find the electronegativies of the two atoms in the bond

Find the absolute value of the difference of their valuesIf the difference is 0.4 or less, it’s a non-polar

covalent bondIf the difference is greater than 0.4 but less than

1.4, it’s a polar covalent bondIf the difference is greater than 1.4, it’s an ionic

bond

Let’s Practice

Example:If the bond

is polar, draw the polarity arrow

C – H

O—Cl

F—F

C—Cl

Polar Bonds versus Polar Molecules

Not every molecule with a polar bond is polar itselfIf the polar bonds cancel out then the molecule

is overall non-polar.

The polar bonds cancel out.No net dipole

The polar bonds do not cancel out.

Net dipole

The Importance of VSEPR

You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not!

Water drawn this way shows all the polar bonds canceling out. But water drawn in

the correct VSEPR structure, bent, shows the polar bonds don’t cancel out!

Net dipole

H O H

O H H

Let’s Practice

Example:Is NH3 a

polar molecule?

Section 5.5—Intermolecular Forces

Intra- versus Inter-molecular Forces

So far this chapter has been discussing intramolecular forcesIntramolecular forces = forces within the

molecule (chemical bonds)

Now let’s talk about intermolecular forcesIntermolecular forces = forces between

separate molecules

Breaking Intramolecular forces

Breaking of intramolecular forces (within the molecule) is a chemical change2 H2 + O2 2 H2O

Bonds are broken within the molecules and new bonds are formed to form new molecules

Breaking Intermolecular forces

Breaking of intermolecular forces (between separate molecules) is a physical changeBreaking glass is breaking the intermolecular

connections between the glass molecules to separate it into multiple pieces.

Boiling water is breaking the intermolecular forces in liquid water to allow the molecules to separate and be individual gas molecules.

London Dispersion Forces

All molecules have electrons.

Electrons move around the nuclei. They could momentarily all “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.

+

-

London Dispersion Forces

Once the electrons have “ganged up” and created a partial separation of charges, the molecule is now temporarily polar.

The positive area of one temporarily polar molecule can be attracted to the negative area of another molecule.

+ - + -

Strength of London Dispersion Forces

Electrons can gang-up and cause a non-polar molecule to be temporarily polar

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 of the intermolecular forces because molecules can’t form it all the time.

Strength of London Dispersion Forces

Larger molecules have more electrons

The more electrons that gang-up, the larger the partial negative charge.

The larger the molecule, the stronger the London Dispersion Forces

Larger molecules have stronger London Dispersion Forces than smaller molecules.

All molecules have electrons…all molecules can have London Dispersion Forces

Dipole Forces

Polar molecules have permanent partial separation of charge.

The positive area of one polar molecule can be attracted to the negative area of another molecule.

+ - + -

Strength of Dipole Forces

Polar molecules always have a partial separation of charge.

Polar molecules always have the ability to form attractions with opposite charges

Dipole forces are stronger than London Dispersion Forces

Hydrogen Bonding

Hydrogen has 1 proton and 1 electron.There are no “inner” electrons. It bonds with the only

one it has.When that electron is shared unevenly (a polar

bond) with another atom, the electron is farther from the hydrogen proton than usual.This happens when Hydrogen bonds with Nitrogen,

Oxygen or FluorineThis creates a very strong dipole (separation of

charges) since there’s no other electrons around the hydrogen proton to counter-act the proton’s positive charge.

Strength of Hydrogen Bond

Hydrogen has no inner electrons to counter-act the proton’s charge

It’s an extreme example of polar bonding with the hydrogen having a large positive charge.

This very positively-charged hydrogen is highly attracted to a lone pair of electrons on another atom.

This is the strongest of all the intermolecular forces.

Hydrogen Bond

N

H H

N

H H

Hydrogen bond

Section 5.6—Intermolecular Forces & Properties

IMF’s and Properties

IMF’s are Intermolecular ForcesLondon Dispersion ForcesDipole interactionsHydrogen bonding

The number and strength of the intermolecular forces affect the properties of the substance.

It takes energy to break IMF’sEnergy is released when new IMF’s are

formed

IMF’s and Changes in State

Some IMF’s are broken to go from solid liquid. All the rest are broken to go from liquid gas.

Breaking IMF’s requires energy.

The stronger the IMF’s, the more energy is required to melt, evaporate or boil.

The stronger the IMF’s are, the higher the melting and boiling point

Water

Water is a very small moleculeIn general small molecules have low melting and

boiling pointsBased on it’s size, water should be a gas under

normal conditionsHowever, because water is polar and can form

dipole interactions and hydrogen bonding, it’s melting point is much higher

This is very important because we need liquid water to exist!

IMF’s and Viscosity

Viscosity is the resistance to flowMolasses is much more viscous than

water

Larger molecules and molecules with high IMF’s become inter-twined and “stick” together more

The more the molecules “stick” together, the higher the viscosity

Solubility

In order from something to be dissolved, the solute and solvent must break the IMF’s they form within itself

They must then form new IMF’s with each other

Solubility

- +

- +

- + - +- +

Solvent, water (polar)

+

-

- + Solute, sugar (polar)

Water particles break some intermolecular forces with other water molecules (to allow them to spread out) and begin to form new ones with the sugar molecules.

Solubility

Solvent, water (polar)

+

-

- + Solute, sugar (polar)

As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving”

- +

- +

- +- + - +

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 occurAn exception to this is if more energy is added

somehow (such as heating)

Oil & Water

Water has London Dispersion, Dipole and hydrogen bonding. That takes a lot of energy to break

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 is polar and can hydrogen bond, Oil is non-polar.

Therefore, oil and water don’t mix!

Surface Tension

Surface tension is the resistance of a liquid to spread out.This is seen with water on a freshly waxed car

The higher the IMF’s in the liquid, the more the molecules “stick” together.

The more the molecules “stick” together, the less they want to spread out.

The higher the IMF’s, the higher the surface tension.

Soap & Water

Soap has a polar head with a non-polar tail

The polar portion can interact with water (polar) and the non-polar portion can interact with the dirt and grease (non-polar).

Polar head

Non-polar tailSoap

Soap & Water

The soap surrounds the “dirt” and the outside of the this Micelle can interact with the water.

The water now doesn’t “see” the non-polar dirt.

Dirt

Soap & Surface Tension

The soap disturbs the water molecules’ ability to form IMF’s and “stick” together.

This means that the surface tension of water is lower when soap is added.

The lower surface tension allows the water to spread over the dirty dishes.

What did you learn about soap?

Soap



Works based on

Molecular Geometry

Molecular Geometry

Bonding types &

Structures

Bonding types &

Structures

Determined by

Determined by

chapter 5: soap. introductory activity fill a test tube with an inch of water add a squirt of...

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