The attractive forces between molecules.

Usually weaker than Intramolecular Forces

(Hint: Interstates are highways that connect other states, intrastates are highways within a state.)

The forces that hold atoms together within a molecule

Usally stronger than intermolecular forces.

(Hint: Interstates are highways that connect other states, intrastates are highways within a state.)

Electronegativity Definition

The ability of an atom in a

molecule to attract the shared

electrons in a covelant bond.

Electromagnetivy Chart.

Increases from left to right.

Increases from top to bottom.

Noble Gases are not Electromagnetic.

Polar Covalent Bond

The bonding electrons are attracted more strongly towards the element that has higher electromagnetity.

Non-Polar Covelant Bond

Electrons are shared equally. The individual bond

polarities cancel. Therefore, the molecule does not

have a dipole moment. In other words, the

molecule is nonpolar.

Bonds by Electromagnetivity

1. Non-Polar Covalent  ?-?

2. Polar Covalent  ?-?

3. Ionic  ?-?

1. Non-Polar Covalent 0 - 0.4

2. Polar Covalent 0.5 - 2.0

3. Ionic 2.0 +

[image]

Dipole Moment

a measure of the net molecular polarity.

The individual bond polarities do not cancel.

Has a dipole moment.

The molecule is polar.

Types of Intermolecular Forces

(The attraction between molecules

that hold them together)

Intermolecular Forces

1. Ion-Dipole Forces

Van der Waals Forces

2. Dipole-Dipole Forces

3. London Dispersion Forces

4. Hydrogen Bonds

Dipole-Dipole Forces

The result of electrical interactions between dipoles on neighboring molecules.

(As the dipole forces increase

the intermolecular forces increase.

As the intermolecular forces increase,

the boiling points increase.)

An attractive force between a hydrogen

atom bonded to a very electronegative

atom (O, N, or F) and an unshared electron

pair on another electronegative atom.

(The strongest of the Van der Waals Forces.

It is a Dipole-Dipole, Dispersion, and Hydrogen Bond.)

London Dispersion Forces Defintion

(Dispersion Forces)

The result of motion of electrons which

gives the molecule a short-lived dipole

momentwhich induces temporary dipoles

in neighboring molecules.

(Temporary Dipoles cause random movement.

All molecules have Dispersion.

It is the weakest of all forces.

Tend to have low boiling points.)

London Dispersion Forces:

The result of motion of electrons which gives the molecule a short-lived dipole moment which induces temporary dipoles in neighboring molecules.

Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted

Polarizability increases with:

• greater number of electrons
• more diffuse electron cloud

As the dispersion forces increase, the intermolecular forces increase.

As the intermolecular forces increase, the boiling points increase.

London Dispersion Forces

Large vs Compact Molecules

Ion-Dipole Forces

The result of electrical interactions between anion and the partial charges on a polar molecule.

(example: salt that can be disolved in a polar molecule.)

Surface Tension Definition

Surface Tension

The amount of engery required to strech or increase the surface of a liquid by a unit of area.

(A physical property.

Strong intermolecular forces = High surface tension)

Viscosity: the measure of a fluid's resistance to flow.

(Strong Intermolecular Forces = High Viscosity)

Phase Change Definition

(State Change)

Phase Change(State Change):

A change in physical form, but not the chemical identity of a substance.

(example: ice melting, water evaporating)

Endothermic Definition

Endothermic reaction:

  A chemical reaction that absorbs heat from its environment.

Positive Enthalpy (+ΔH) and Positive Entropy (+ΔS)

Exothermic:

A chemical reaction that releases heat and has a negative enthalpy (-ΔH) and negative entropy (-ΔS).

Fusion (melting):

From a solid to a liquid

Endothermic reaction-absorbs heat

Positive Enthalpy (+ΔH) and Positive Entropy (+ΔS)

(example: melting ice or solid salt)

Vaporization Definition

Vaporization:

From a liquid to a gas

Endothermic reaction-absorbs heat

Positive Enthalpy (+ΔH) and Positive Entropy (+ΔS)

(example: water to vapor)

Sublimation:

From a solid to a gas

Endothermic reaction-absorbs heat

Positive Enthalpy (+ΔH) and Positive Entropy (+ΔS)

(example: dry ice)

Freezing:

From a liquid to a solid

Exothermic reaction-releases heat

Negative Enthalpy (-ΔH) and Negative Entropy (-ΔS)

(example: water to ice)

Condensation:

From a gas to a liquid

Exothermic reaction-releases heat

Negative Enthalpy (-ΔH) and Negative Entropy (-ΔS)

(example- water vapor becomes liquid on a glass,

dew, clouds)

Depostion:

From a gas to a solid

Exothermic reaction-releases heat

Negative Enthalpy (-ΔH) and Negative Entropy (-ΔS)

(example: water vapor freezing into frost in winter)

Enthalpy of (Heat) Fusion Definition

(ΔH fusion)

Enthalpy of (Heat) Fusion (ΔH fusion):

The amount of energy required to convert a solid into a liquid.

Enthalpy (Heat) of Vaporization Definition

(ΔH vapor)

Enthalpy (Heat) of Vaporization (ΔH vapor):

The amount of energy required to convert a liquid into a gas.

Vapor Pressure:

The partial pressure of a gas in equilibrium with liquid at a constant temperature.

Temp ↑ = vapor pressure ↑

Boiling Point:

The temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure.

Normal Boiling Point:

The temperature at which a liquid boils when the external pressure is 1 atm,

or 760mm of Hg.

Phase Diagram:

Summarizes the conditions at

which a substance exists as a

solid, liquid, or gas.

Normal freezing Point:

Normal Boiling Point:

Normal freezing Point:

Normal Boiling Point:

both occur at 1 atm or 760mm of Hg

Critial Temperature Definition

 +   Critial Pressure Definition

= Critial Point Definition

Critial Temperature:  the temperature beyond which a gas can be liquified reguardless of pressure.

Critial Pressure:   the pressure beyond which a gas can be liquified reguardless of temperature.

Critial Point:      Combination of temperature and pressure beyond which a gas can be liquified.

Supercritical Fluid:

A state of matter beyond the critial point that is neither liquid nor gas.

Triple Point:

The temperature and pressure at which point all three phases coexist in equilibrium.

(Solid, liquid, and gas)

(example: water at 0.01°C)