Chapter 5: Liquids and Solids

The attractive and repulsive forces between molecules in liquids and solids are some of the most important forces in Chemistry– Why different liquids boil at different temperatures– Why some liquids flow and others don’t– Why diamonds are the hardest known substances– Why DNA twists into a double helix

• In this chapter, we’ll study Intermolecular Forces: the forces that exist between molecules

Phases

• Last chapter we studied gases• According to one of the tenets of the kinetic

molecular theory– The molecules do not attract each other

• In liquids and solids, because of the proximity of the molecules to each other, the intramolecular forces between molecules ARE significant

• We’ll refer to liquids and solids as Condensed phases

Interion and Intermolecular Forces

•Ion-Ion interactions are the strongest interactions

•Example of an ion-ion interaction?

•Let’s look at the various interactions given in the table

Ion-Dipole Interactions

•Best example: Hydrated Ions

•The polar character of the water molecule allows it to interact with cations or anions

•We can describe the interaction energy:

€

E p ∝ -zμ

r2

z = ion charge

µ = Electric dipole moment

r = distance

Hydrated vs Anhydrous Molecules

• When salts crystallize out of water, they sometimes pull water molecules with them

Na2CO3•10H2O

CuSO4•5H2O

• These are called hydrated salts• The Group I and II metal cations form much

more hydrated salts than transition metal salts– These cations are smaller (usually) than transition

metal cations

Remember: Molar mass calculation includes the water molecules!

Ion-dipole interactions are strong for small, highly charged ions; one consequence is that small, highly charged cations are often hydrated in compounds

Dipole-Dipole interactions

• Let’s look at the interactions between polar molecules

The Dipole-Dipole interactions force some order in the solution

Dipole-Dipole interactions

Dipole-Dipole interaction energy:

€

E p = -μ1μ2

r3

µ1: Dipole moment of molecule 1

µ2: Dipole moment of Molecule 2

The fact that the distance is cubed means that the energy falls of much more rapidly than ion-ion interactions as the interacting species are separated

Which molecule has the higher boiling point:

p-dichlorobenzene and o-dichlorobenzene

Dipole moment for the molecules?

Which molecule has the higher boiling point:

cis-dichloroethene or trans-dichloroethene?

London Forces

• London Forces are attractive forces between non-polar molecules

• These are 1 of the two weakest intermolecular forces

• How do these interactions arise?

London Forces

• The electron clouds are constantly shifting and sometimes the molecule gets a small dipole moment– Neighboring nonpolar molecules will

have their electron clouds distorted and will form a dipole of opposite orientation

• Then the process starts over (Dipole disappears and reforms) (1x10-16 sec to form and disappear)

London Forces

€

E p = -α 1α 2

r6

1: Polarizability of molecule 1

2: Polarizability of molecule 2

r6 !!!! Very short range effects!!

•What determines Polarizability?

•Large atomic radii

•Low Zeff

•High Polarizability = Large London Interactions

Let’s look at London Forces and Polarizability with respect to physical properties

As we go down a group, the atomic radius increases and the melting and Boiling points increase (takes More energy)

London Forces and Molecular Shape• Because the London Force energy drops off

VERY sharply as a function of distance, molecular shape is a major contributor to London Force energy

Which has the higher boiling point?

Dipole - Induced Dipole

• The presence of a molecule with a strong dipole moment can induce or create a dipole in a non-polar molecule– This depends on the strength of the dipole

and the polarizability of the nonpolar molecule

€

E p = -μ1α 2

r6

1: Dipole moment of molecule 1

2: Polarizability of molecule 2

Hydrogen Bonding

• A special type of dipole-dipole interaction

• Hydrogen bonding only occurs between:

N-H

O-H and Lone pair e- on N, O, F

F-H

Hydrogen Bonding

• Hydrogen bonds are one of the most important interactions in biological systems

Hydrogen Bonds:

•Hold proteins together

•Allow DNA base pairs to match up

•Allow structural polymers to interact

Hydrogen bonds are the strongest type of intermolecular force

Liquids

• The liquid phase is a unique phase midway between the solid phase and the gas phase

• In the liquid phase, the molecules move past and around one another, but they don’t reach the separation distances achieved by molecules in the gas phase

– There is not enough energy to overcome all of the intermolecular forces present in the system

Liquids and Short Range Order

• In the liquid state, we can look at a molecule and its immediate neighbors and see the interactions occurring

• However, we cannot say what the exact arrangement will be some distance away from where we are looking– This is an example of Short Range Order

Viscosity

• You’ve heard of viscosity before, but probably only in commercials on television for motor oil

• Viscosity is defined as the resistance of a liquid to flow

• Viscosity is dependent on temperature to some extent– As we pump more thermal energy into the system,

the molecules can break their intermolecular interactions and move about more easily

• Viscosity Decreases as Temperature Increases

Viscosity: Examples

• Oils: Carbon chains / London Forces

• Water: Hydrogen Bonding

• Sulfur: S8 ring breakage

Surface Tension

• Why is the surface of a liquid smooth?

• The intermolecular forces in the liquid

• The tendency for liquids to form droplets is a direct result of surface tension– Water forms droplets– Nonpolar molecules like

chloroform only have London Forces and do not form as large a droplet

“Wetting”

• Water is “wet” because it can form hydrogen bonds with a surface and it does so with ease

• Let’s think about this in terms of things you’ve seen before– Water on a waxy leaf: It beads up; It doesn’t “wet”

the leaf– Water on Paper: The water “wets” the paper

because the cellulose polymers in the paper have lots of -OH groups

Capillary Action

• Capillary action is the rise of liquids up a narrow tube

• Happens as a result of the sum of the adhesive forces and cohesive forces of the molecule and the glass

• Adhesion: Forces/Interactions between a liquid and a surface

• Cohesion: Intermolecular forces that keep molecules bound together

Solids

• Solids are formed when temperatures are low enough to prevent molecules from moving around their neighbors

• Molecules or atoms in solids vibrate/oscillate around a fixed position– All the atoms in a solid essentially move in

unison

Types of Solids

• Two major classifications: Crystalline and Amorphous

1. Crystalline Solids: All atoms, ions or molecules in the solid lie in an orderly array. Long range order exists within a crystalline solid

2. Amorphous Solids: The atoms, ions or molecules in an amorphous solid lie in a seemingly random arrangement.

• Amorphous solids look like we took a snapshot of a liquid

• Glass is an amorphous solid

Crystalline Solids

1. Molecular Solids: Assemblies of discrete molecules held together by intermolecular forces (eg: Quartz)

2. Network Solids: Atoms are covalently bonded to their neighbors throughout the extent of the solid (eg: Diamonds)

3. Metallic Solids: Metal cations held together by a sea of electrons

4. Ionic Solids: Built from the mutual attraction of anions and cations