Chapter 12 – Liquids and Solids

1.  The three states of matter:  (there are actually 2 others: Plasma and the Bose-Einstein Condensate but we will consider only the basic three which we commonly find on Earth)

            Solid:  A substance with a definite volume and shape – most highly organized

            Liquid:  A substance which flows and takes the shape of its container – less organized than a solid.

            Gas:  A substance which takes the shape of and fills its container – less organized than a liquid.

            Kinetic energy:  Energy of an object because of its motion

2. Changing States of matter:  See table 12-2 p. 372 and diagram 12-15 on p. 382

            Intermolecular forces:  The force of attraction between two or more molecules.  Gases have the weakest/lowest intermolecular forces, solids have the strongest/largest.

            Condensation:  This is the process of going from a gas to a liquid.  If the temperature is lowered enough, the average kinetic energy of a particle is too low to remain as a gas, therefore, it is attracted to other gas particles and condenses into a liquid.  If intermolecular forces are large, then the condensation will happen at a higher temperature.  Example:  Water has higher intermolecular forces than alcohol.  Water condenses from a vapor to a liquid at 100 degrees while alcohol might do it at 80 degrees.  This is because it takes more energy to break water apart from liquid to gas.

            Boiling or Vaporization or Evaporation: (p. 378)  This is the process of going from a liquid to a gas.  It is the exact opposite of condensation.  The higher the intermolecular forces, the higher the evaporation temperature.  Boiling temp. and Condensation temp. are exactly the same.

You can alter the boiling/vaporization/evaporation point of a liquid by altering the pressure on the liquid.  Polar compounds will have a higher boiling point than non-polar compounds.

            Freezing (p. 379-380):  This is the process of a liquid becoming a solid.  The higher the intermolecular forces, the higher the freezing point.  For Example:  The metal pipes in the room have very strong intermolecular forces.  Their freezing point is very high – several hundred degrees.  You can alter the freezing point by altering the outside pressure.  Polar compounds will have a higher freezing point than non-polar compounds.

     Look at table 12-1 on p. 370 and you will see that covalent molecules (which have lower intermolecular forces than ionic compounds) have lower melting and boiling points.

            Melting or Fusion (p. 379-380):  This is the process of a solid becoming a liquid.  It happens at the exact same temperature as the freezing point.

            Sublimation (p. 380):  This is the process of a solid becoming a gas and NEVER entering the liquid state.  The reverse of this is called Deposition. 

            Note:  Refer to the phase diagrams on p. 381 (water) and 390 (CO2) to help you see where these phases of matter changes occur as a function of pressure and temperature.

3.  Let’s take a trip down memory lane…….:

     Remember from chapter 6:  Intramolecular forces are the forces which hold individual atoms together to make a molecule.  These forces are determined by the atoms’ individual electronegativity rating (see p. 143).  Intermolecular forces, on the other hand, are the forces between molecules holding them together.  For example, water molecules clinging together in a glass of water.  You will need to be able to distinguish between inter and intra molecular forces.

 

4.  Equilibrium (p. 372-375):  Equilibrium is a dynamic (ever changing) condition in which two opposing changes occur at equal rates in a closed system. Example:  Stand on one leg and close both of your eyes – you will lose your equilibrium of balance.  Example:  Look at the Sparkletts water bottle in class.  Notice that it is in equilibrium between liquid and gas phase.

     Le Chatelier’s Principle:  When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position which minimizes the stress. (p. 374)  See table 12-3 on p. 375 and understand what changes in conditions will cause a shift.

 

5.  Vapor Pressure:   (page 376 figure 12-12) This is the pressure of the vapor above a liquid at a given temperature.  Example:  As water is heated to 100oC, its vapor pressure approaches that of the air (1 atm).  As Ethanol is heated to 80oC, its vapor pressure approaches that of the air.   This figure shows that water must have the highest INTERMOLECULAR forces among these three molecules.  It takes a lot of heat to make water reach 1 atm at its surface.   When vapor pressure reaches atmospheric pressure (1 atm or 101.325 kPa), the liquid boils. 

     Strong Intermolecular forces make for LOW vapor pressures.  Water, for example has strong intermolecular forces.  Its vapor pressure is very low.  This means that not much of the water is in vapor form above the liquid.  Water wants to remain a liquid and therefore there is not much vapor above a swimming pool (for example).  You must heat the water up a lot in order to increase the vapor pressure to 1 atm.

    

     Bottom line:  If a liquid has a low vapor pressure, it is going to have a high boiling point.  If a liquid has a high vapor pressure, it is going to have a low boiling point – because its liquid is readily available to converting to a gas.   (Make sure you read p. 374-377 to re-enforce this idea)

 

            Volatile liquids:  These have low intermolecular forces and very high vapor pressures.  They are ready to convert to a gas at a moments notice.  Gasoline would be an example of a volatile liquid with a high vapor pressure and a low boiling point.  Water is not volatile.

            How to manipulate boiling point:  Boiling point is the temperature at which the vapor pressure equals the external pressure acting on a liquid.  If you reduce the external pressure on the liquid, then the vapor pressure will reach that external pressure at a lower temperature.  The same is true for water.

            How to manipulate freezing point:  Reduce the outside pressure and the freezing point will decrease.  Without outside pressure pushing the molecules together, the molecules will have to be COLDER to fuse together to make a solid. EXCEPT FOR WATER!  With water, as pressure increases, the freezing point will decrease (look at phase diagram on p. 381)  This is why, in the mountains, where there is less pressure, it can snow more easily.  It doesn’t have to be as cold in the mountains (as it does as in San Clemente) to make snow fall.

   Compare the phase diagram of water with that of Carbon dioxide on p. 390

 

6.  Water’s properties

            Density of water based on temperature:  See page 385 figure 12-17.  As water freezes to form ice, the hydrogen bonding which occurs allows for spaces to open up between the molecules of water and therefore water is less dense as a solid than as a liquid.

     As water freezes, it expands.  This is why bottles blow up in the freezer and pipes may burst in the winter in cold climates.

     Water is at its most dense right before it freezes.  At 4 degrees Celsius, water is at its most dense and it sinks to the bottom of a lake (for example).  But at 0 degrees, it is at its least dense and therefore it rises.  (Realize that 4 degree water is a liquid but 0 degree water is a solid)

            Water’s heat capacity:  Because of the strong hydrogen bonding, water has a high heat capacity.  It takes a lot of heat to warm up water.  And water can store a lot of heat too.  Substances with low heat capacity (such as metals) heat up quickly and then cool down quickly too.

 

 

 

           

Chapter 12 – Liquids and Solids

Study Guide

 

You must be able to define the following terms and know how they are related.

 

1.  The four states of matter: (look in the glossary in the back of the book)

            Solid

            Liquid

            Gas

           

 

2. Changing States of matter:

            Intermolecular forces

            Condensation

            Vaporization or Evaporation

            Boiling

            Freezing

            Melting or Fusion

            Sublimation

            Phase diagram of water

            Phase diagram of CO2.

 

3.  Intermolecular Forces holding molecules together: (Review Chapter 6)

            Electronegativity

            The van der Waals forces:

                        Dispersion forces

                        Dipole-Dipole forces

                        Hydrogen Bonding

 

4.   Equilibrium

              Le Chatelier’s Principle

 

5.  Vapor Pressure

            Partial pressure

            Volatile liquids

            How to manipulate boiling point

            How to manipulate freezing point

 

6.  Water’s properties

            Density of water based on temperature

            Water’s heat capacity

 

7.  Common Temperatures of substances at 1 atm of pressure

                                                 oC             oF                                          

Human Body                           37             98.6

Freezing Water                         0               32

Boiling Water                         100            212

Dry Ice (CO2)                         -78             -94

Liquid Nitrogen (N2)             -196           -321

Room Temperature                  25              77