Chemistry 103A; Sections 5, 6, 7, 8; Lecture 40, 1 Dec 00

We were talking about:

Molecular and Network Solids

Molecular Solids

Ice, dry ice, iodine are examples of molecular solids. In molecular solids the smallest unit of the compound is the molecule. In the solid phase the molecules pack themselves together in the way that give the ensemble the minimum energy.

The forces that hold molecular crystals together are the familiar intermolecular forces, dispersion forces, dipole-dipole forces, and hydrogen bonding forces.

(Hydrogen bonding forces are responsible for the unusual property of water that the solid is less dense than the liquid at 0 oC. When water freezes the average distance between water molecules increase slightly to allow room for the formation of hydrogen bonds. Since hydrogen bonding in the dominant intermolecular force in water, forming more hydrogen bonds will lower the energy of the crystal.) Network Solids

The most familiar network solids are diamond and graphite. Many minerals are network solids. For example, quartz (SiO2), silicates, silicon carbide, etc. Boron nitride, BN, has the same crystal structure as diamond and it is reputed to be slightly harder than diamond.

A network solid, such as a diamond or graphite crystal, is essentially one giant molecule. The forces that hold the solid together are covalent bonds. One could start with an atom on one side of a macroscopic crystal and trace a path all the way to the other side of the crystal moving from atom to atom along covalent bonds.
 

Solutions

We talked a little bit about solutions in Lecture 8 (Chapter 5) Solutions are members of the class of substances we called mixtures. Recall that there were two categories of mixtures:

homogeneous mixtures and

inhomogeneous mixtures.

Solutions are homogeneous mixtures.

Add sugar or salt to water - stir - the sugar or salt seems to disappear.

We say the sugar or salt dissolved to form a solution.

A binary solution is a solution that contains two components. We can talk about the two components as:

solvent = the major component, and

solute = the minor component.

In a solution the solute is separated into individual molecules, atoms, or ions, and distributed evenly throughout the solvent.

Main characteristics of a solution

1) Homogeneous

2) Stable

3) Can't be separated by filtration

4) Continuously variable composition

5) Usually transparent

6) Can be separated (distillation, etc)

Types of solutions

Solvent              Solute             Examples

gas                       gas                      air

(We usually don't think of gas mixtures as solutions, but they satisfy all the criteria.)

liquid                    gas             carbonated water

liquid                 liquid          vodka, engine coolant

liquid                  solid                 sea water

(champagne is all three)

Solvent             Solute            Examples

solid                    gas             H2 in Pt or Ir

solid                  solid                  alloys

In this chapter we will discuss mainly solutions were the solvent is a liquid.
 

Concentration Units

The concentration is a measure of the relative amounts of solute and solvent. There are many units of concentration, of which the most important for us are:

Molarity (We have already used molarity in Chapter 8.) Molarity is defined by,  
.


(We stress once again that molarity is mol of solute per Liter of solution, not per liter of solvent. Molarity is defined so that we can always know how many mols of solute there are in any given amount of solution.)

Molality

Molality is similar to but not the same as molarity. Molality, m, is defined by,

 

Note that for water solutions 1.00 kg of water has a volume of 1.00 L. If the solution is dilute the volume of the solution formed from 1 L of water is still approximately 1 L, so that the molarity and molality are about the same. However, in concentrated water solutions and in solutions where the solvent is not water the molarity and molality are very different.


Mole Fraction

 
Mole fraction is essentially self-defined. In equation form the mole fraction (usually symbolized by X) is  
Other units which are important in medicine and medical technology, food, and in monitoring the environment are:
Weight percent (wt%)
Weight percent is defined as,
Volume percent or percent by volume (vol%)  
Percent by volume is usually used for liquid-liquid solutions, as in alcoholic beverages. It is defined by,  

(The "proof" of an alcoholic beverage is twice the vol% concentration of ethyl alcohol. A 100 proof beverage is 50 vol% alcohol.)

 wt/vol%  
Perhaps the most common of the "%" type concentration units is the wt/vol%. This unit is defined by,  
Note that the units in this one matter. The units of wt/vol% are g/100mL.


Parts per million (ppm)

 
Parts per million is usually used to describe the concentration of trace contaminants in otherwise pure materials. ppm is defined by  


Since for very dilute solutions the mass of the solution and the mass of the solvent are very nearly the same one could just as well write,


Parts per billion (ppb)

 
Nowadays, with increasingly accurate methods of analysis and increasing concern over minute amounts of some suspected pollutants, we frequently see reports with concentrations in the ppb range. ppb is defined similarly to ppm as,  
Example calculation:

Calculate the molarity of a 0.89 wt% aqueous solution of NaCl.
 
 

Solubility

The solubility of a substance (solute) is the amount of solute that will dissolve in a given amount of solvent. It is usually expressed in units of g solute/100mL solvent. The solubility is a measured quantity and there are tables of solubility for various substances.

Solubility depends on temperature. For most solutes the solubility increases with increasing temperature. A solution is said to be unsaturated if the amount of solute is less than the solubility. The solution is saturated if the amount of solute in solution is equal to the solubility. It is possible to prepare a supersaturated solution in which the amount of solute in solution is greater than the solubility. This is not an equilibrium state and a disturbance of the solution will cause the excess solute to precipitate out. (We say that a super saturated solution is metastable in the same way that super heated and super cooled liquids are metastable.)

A supersaturated solution can be prepared by dissolving as much solute as possible in hot solvent and allowing the solution to cool slowly and quietly to room temperature. Since there are random probabilities involved this will not always work, but sometimes it will work.
 

 Properties of Solutions

Solubility of Gases in liquids

The solubility of a gas in a liquid is proportional to the pressure of the gas. The mathematical expression of this phenomenon is called "Henry's law." Henry's law has the form,

where Sg is the concentration of the gas in the solution (in mol/L) and pg is the pressure of the gas above the solution (in Torr). The constant, kH, is the proportionality constant and is called the Henry's law constant. The Henry's law constant depends on the solvent, the solute, and the temperature. These constants are measured and listed in data tables. There is at table of Henry's law constants for several gases in water on p. 653 in your text.

Example,

What is the concentration of oxygen in water open to the atmosphere at 25oC? (The partial pressure of O2 in air with a total pressure of one atmosphere is approximately 0.20 atm = 150 Torr. kH for oxygen in water at 25oC is 1.66 ´ 10-6 mol/Ltorr.)

Using Henry's law we find

which is about 8 mg per liter.

 The solubility of gases in liquids usually decreases with increasing temperature unless there is a reaction between the gas and the liquid (such as HCl and water). Boiling a liquid will generally drive the gases out of it. You would not want to use boiled water in your fish tank unless you first aerated it well.