ch 11 states of matter and intermolecular forces

Ch 11 States of Matter and Intermolecular Forces

Intramolecular forces (bonds) govern molecular properties.

Intermolecular forces are between molecules and determine the macroscopic physical properties of liquids and solids.

This chapter: describes changes from one state of matter to

another. explores the types of intermolecular forces that

underlie these and other physical properties of substances.

Chapter 11 Preview

Review:Intramolecular Forces

Ionic Covalent Metallic

Ionic bond- electron donated and accepted

Covalent- sharing electrons

Metallic-sea of electrons

Ionic Bonds as“Intermolecular” Forces

There are no molecules in an ionic solid, and therefore there can’t be any intermolecular forces.

These forces increase: as the charges on the ions increase. as the ionic radii (size) decrease.

Interionic Forces of Attraction

Melting point of NaCl is about 801 oC.

Mg2+ and O2– have much stronger forces of attraction for one another than do Na+ and Cl–. Melting point of MgO is about 2800 oC.

Molecular Forces Compared

States of Matter ComparedIntermolecular

forces are of little significance; why?

Intermolecular forces must be

considered.

Intermolecular forces are very

important.

Intermolecular Forces

Hydrogen bonding Dipole-dipole London Forces

Cohesion

Attraction for each other Water Mercury

Boiling point varies based on cohesion

Adhesion

A liquids attraction for solid particles Water’s attraction for glass etc.

WATER MERCURY

Meniscus Formation

Water wets the glass (adhesive forces) and its

attraction for glass forms a concave-

up surface.

What conclusion can we draw about the cohesive forces

in mercury?

Plant root- capillary action

Surface Tension The forces present in liquids A molecule inside the liquid experiences

cohesive forces with other molecules in all directions.

A molecule at the surface experiences only downward cohesive forces

Surface TensionThere is no force above the molecule on the top of the

surface.

Pool ball floating on mercury

Adhesive and Cohesive Forces

The liquid spreads, because adhesive forces

are comparable in strength to cohesive

forces.

The liquid “beads up.” Which forces are

stronger, adhesive or cohesive?

Cohesive vs. Adhesive Water on plastic Water on metal Water on glass

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Intermolecular Forces

Hydrogen bonding Dipole-dipole London Forces

Hydrogen bonding A hydrogen is attracted to a highly

electronegative atom like O, N, or Cl.

Hydrogen Bonding in Water

Hydrogen Bonding in Ice

Hydrogen bonding arranges the water molecules into an

open hexagonal pattern.

“Hexagonal” is reflected in the crystal structure. “Open”

means reduced density of the solid (vs. liquid).

Hydrogen Bonding in Acetic Acid

Hydrogen bonding occurs between

molecules.

Intermolecular Hydrogen Bonds

Intermolecular hydrogen bonds give proteins their secondary shape, forcing

the protein molecules into particular orientations, like a folded sheet …

Intramolecular Hydrogen Bonds

… while intramolecular

hydrogen bonds can cause proteins to

take a helical shape.

In which of these substances is hydrogen bonding an important intermolecular force: N2, HI, HF, CH3CHO, and CH3OH? Explain.

In which of these substances is hydrogen bonding an important intermolecular force: N2, HI, CH3CHO, and CH3OH? Explain.

CH3CHO and CH3OH because of the attraction between the H of one molecule and the O of another.

HI would not have hydrogen bonding b/c iodine is not highly electronegative.

Dipole–Dipole Forces

A polar molecule has a positively charged “end” (δ+) and a negatively charged “end” (δ–).

When molecules come close to one another, repulsions occur between like-charged regions of dipoles. Opposite charges tend to attract one another.

Dipole Forces

The more polar a molecule, the more pronounced is the effect of dipole–dipole forces on physical properties.

Dipole–Dipole Interactions

Opposites attract!

London Forces-aka Dispersion Forces

At first no dipole, like Argon. But the electrons are mobile, and at any one

instant they might find themselves towards one end of the molecule, making that end -. The other end will be temporarily short of electrons and so becomes +.

An instant later the electrons may have moved up to the other end, reversing the polarity of the molecule.

Induced London ForcesWhat would happen if we mixed HCl with the

element argon, which has no dipole? The electrons on an argon atom are distributed

homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.

Dispersion Forces Illustrated (1)At a given instant, electron density, even in a nonpolar

molecule like this one, is not perfectly uniform.

Dispersion Forces Illustrated (2)

The region of (momentary) higher

electron density attains a small (–) charge …

When another nonpolar molecule

approaches …

… the other end of the molecule is

slightly (+).

Dispersion Forces Illustrated (3)… this molecule induces a tiny

dipole moment …

… in this molecule.

Opposite charges ________.

Types of forces

Click here for an Animation of the forces

Molecular Shape and PolarizabilityLong skinny molecule …

… can have greater separation of charge along its length. Stronger forces of attraction, meaning …… higher

boiling point.

In the compact isomer, less possible separation of charge …

… giving weaker dispersion forces and a lower boiling point.

Arrange the following substances in the expected order of increasing boiling point: Carbon tetrabromide, CBr4; Butane, CH3CH2CH2CH3; Fluorine, F2;Acetaldehyde, CH3CHO.

Arrange the following substances in the expected order of increasing boiling point: Carbon tetrabromide, CBr4; Butane, CH3CH2CH2CH3; Fluorine, F2;Acetaldehyde, CH3CHO.

Answer: F2, CBr4, CH3CH2CH2CH3, CH3CHO

Vapor Pressure

The vapor pressure of a liquid is the partial pressure exerted by the vapor when it is in dynamic equilibrium with the liquid at a constant temperature.

vaporizationLiquid Vapor

condensation

Liquid–Vapor Equilibrium

More vapor forms; rate of condensation of that

vapor increases …

… until equilibrium is attained.

Phase DiagramsA phase diagram is a graphical representation of the conditions of temperature and pressure under which a substance exists as a solid, liquid, a gas, or some combination of these in equilibrium.

A—B, solid-vapor equilibrium.

A—D, solid-liquid equilibrium.

A—C, liquid-vapor equilibrium.

Triple point

Phase Diagram for CO2

Note that at 1 atm, only the solid and

vapor phases of CO2 exist.

Phase Diagram for H2O

Supercritical Fluid

Above the critical temperature and pressure, only one phase exists…a combination of liquid and gas.

Properties are in between those of liquids and gases.

They act as solvents and dissolve well. They diffuse well like gases. CO2 and H2O- environmentally friendly

The Critical PointAt room temperature there is relatively little vapor, and

its density is low.

At higher temperature, there is more vapor, and its

density increases …

… while the density of the liquid decreases; molecular

motion increases.

At Tc, the densities of liquid and vapor are equal; a single

phase.