CHM 101 GENERAL CHEMISTRY

FALL QUARTER 2008

Section 2

Lecture Notes – 10/22/2008

(last revised: 10/22/08, 3:45 PM)

4.1 Water, the Common Solvent

• You may recall from Chapter 3, that we discussed labeling each reagent in the equation for a chemical reaction by its state. And that we included aqueous solution in the list of states, along with solid, liquid, and gas. Some of you may have correctly inferred that chemistry taking place in solution is a very important topic. And that among all the possible solvents, water is the most important.
• What Makes Water Such a Good Solvent? You may recall from Chapter 2 that water has a “bent” geometry with the two carbon-hydrogen bonds intersecting at an angel of about 105°. These bonds are covalent, meaning that they involve the sharing of electrons between the oxygen atoms and the hydrogens. However, the sharing is not equal – the electrons in an oxygen – hydrogen bond are more strongly attracted to the oxygen than to the hydrogen. As a result, the atoms in a water molecule are not electrically neutral – The oxygen has a small net negative charge (less than one unit), while the hydrogens each have small net positive charges. This is illustrated in Figure 4.1 from the text. Water is said to be a polar molecule on account of this unequal charge distribution, and it is this polarity that makes water such a good solvent.
• How Ionic Substances Dissolve in Water: Recall that an ionic substance such as the salt, sodium chloride (NaCl) is held together with ionic bonds, i. e., electrostatic attractive forces between oppositely charged ions. Clearly if an ionic solid is to dissolve in water, there must be attractive forces between water and ions of either charge. But as we have just seen, the needed charges do exist in the water molecule. A cation such as sodium ion (Na⁺) will be attracted to the partial negative charge on the oxygen atom of water, and an anion such as chloride (Cl^-) will be attracted to the partial positive charges on the hydrogens. If these attractive forces are more favorable than the forces that hold the ionic solid together, it will break up, and the ions will disperse throughout the water. Each dissolved ion will be surrounded by water molecules.
• How Sodium Chloride Dissolves in Water: The dissolution of sodium chloride in water can be written as the equation:

	H2O
NaCl (s)	——>	Na^+ (aq) + Cl^- (aq)

In the equation, Na⁺ (aq) represents a hydrated sodium ion, or a sodium ion surrounded by water molecules, with the partially negatively charges oxygen atoms oriented inward toward the positively charged sodium ions. Similarly, Cl^- (aq) represents a chloride (ion) surrounded by water molecules with their hydrogens oriented inward. Figure 4.2 from the text gives this pictorial representation of the dissolution process. In the upper panel, we see two water molecules “tugging” at a chloride with their hydrogens and two other water molecules tugging on a sodium ion with their oxygens. The second panel shows two hydrated ions, one each of sodium ion and chloride, and it shows other water molecules tugging at still more ions from the solid lattice:

• How Nonionic Substances Dissolve in Water: Water is not a solvent for all substances. (How fortunate for us?) Many ionic solids, such as silver chloride (AgCl) do not dissolve in water. The forces holding the solid AgCl lattice together are too strong to be overcome by the forces favoring the formation of the hydrated ions, Ag⁺(aq) and Cl^-(aq). Covalently bonded, non polar substances like fats do not dissolve in (pure) water because there is no driving force to form hydrated fat molecules. However, covalently bonded substances like sugars and ethanol will readily dissolve in water. The reason is that ethanol and sugars, like water, are polar because, like water, they contain hydrogen-oxygen bonds. Figure 4.3a illustrates the polarity of the O—H bond in ethanol, and Figure 4.3b pictures the attraction between ethanol and water that promotes solubility:

4.2 Strong and Weak (and non) Electrolytes

• Solvent: We speak of a substance that dissolves something as a solvent.
• Solute: The “something” that is dissolved in a solvent is called the solute.
• Homogeneous Mixture: Recall from Chapter 2 that a solution is homogeneous. The solute is mixed throughout the solution, and the composition of the solution is the same everywhere in the solution.
• Variable Composition: The composition of a solution of A (the solute) in B (the solvent) is not fixed. Two solutions can be prepared with the same amounts of solvent but different amounts of the solute and therefore have different compositions.
• Electrical Conductivity: A useful property of solutions is their electrical conductivity. Pure water does not conduct electricity, but many aqueous solutions do.

• Strong Electrolytes: Strong electrolytes are solutions whose electrical conductivity is high. They are good conductors of electricity.
• Weak Electrolytes: Weak electrolytes are solutions that conduct electricity, but not very well. They are still called conductors, but only poor conductors.
• Non Electrolytes: Non electrolytes do not conduct electricity.

A device for detecting electrical conductivity in solutions is shown in Figure 4.4 in the text:

• Arrhenius’ Theory of Electrical Conductivity: The first modern theory to explain electrical conductivity in solutions was advanced by Svante Arrhenius, at the time a doctoral student in physics at the University of Uppsala, Sweden. He deduced the presence of charged particles in an electrically conductive solution, and he postulated that the extent of the conductivity depended on the number of ions present. Hence the solute in a strong electrolyte would produce many ions and conduct electricity well, while the solute in a weak electrolyte would produce relatively few ions and not conduct electricity very well. In a non electrolyte, the solute would produce no ions, hence no electrical conductivity for these solutions.
• Strong Electrolytes: The conductivity of a strong electrolyte is high, as illustrated in Figure 4.4a. This conductivity is produced by solutes that are completely ionized in solution. Three classes of strong electrolytes are:

o Soluble Salts

o Strong Acids

o Strong Bases

• Soluble Salts: Any salt that readily dissolves in water produces a strong electrolyte when it is so dissolved. A good example is sodium chloride (NaCl). See Figure 4.5.

• Strong Acids: According to Arrhenius, an acid is a substance that ionizes in aqueous solution to generate H⁺ ions (hydrogen ions or protons). If the ionization is complete or nearly complete, the acid solution is a good conductor of electricity, and the acid is regarded as a strong acid. Some examples of strong acids in aqueous solution are:

	H2O
HCl	——>	H^+(aq) + Cl^-(aq)
	H2O
HNO3	——>	H^+(aq) + NO3^-(aq)
	H2O
H2SO4	——>	H^+(aq) + HSO4^-(aq)

Sulfuric acid (H₂SO₄) deserves a closer look. When sulfuric acid dissolves in water, the first hydrogen dissociates completely to form protons in aqueous solution, but the second hydrogen remains bonded to the sulfate. Thus aqueous sulfuric acid contains mostly protons (H⁺(aq)) and hydrogen sulfate (HSO₄^-(aq)) ions.

• Strong Bases: According to Arrhenius, a base is a substance that ionizes in aqueous solution to produce hydroxide ions (OH^-). Bases that ionize completely are regarded as strong bases. The two most common strong bases are sodium hydroxide and potassium hydroxide:

	H2O
NaOH (s)	——>	Na^+(aq) + OH^-(aq)
	H2O
KOH (s)	——>	K^+(aq) + OH^-(aq)

• Weak Electrolytes: Many substances will form ions in aqueous solution, but the extent of ionization is slight. For example, acetic acid (the essence of vinegar), will readily dissolve in water, but only around 1% of acetic acid molecules will ionize to form hydrogen ions and acetate ions. The remaining acetic acid molecules will remain as electrically neutral acetic acid molecules, even though they are totally dissolved in the solution. The result is a solution that is electrically conductive, but much less so than a comparative solution containing a strong electrolyte. The most common weak electrolytes are:

o Weak acids

o Weak bases

• Weak Acids: A weak acid is an acid that is only partially ionized in aqueous solution. Thus an aqueous solution of acetic acid (HC₂H₃O₂) contains some hydrogen ions (H⁺(aq)) and some acetate ions (C₂H₃O₂^-(aq)), but most of the solute particles are undissociated acetic acid molecules (HC₂H₃O₂(aq)).

	H2O
HC2H3O2(aq)	——> <——	H^+(aq) + C2H3O2^-(aq)

• Weak Bases: Ammonia (NH₃) is the most common of the weak bases. It is a base because its aqueous solutions contain hydroxide ions (OH^-(aq)). It is a weak electrolyte because only a small fraction of ammonia molecules form ions. Most of the ammonia remains as neutral ammonia molecules.

	H2O
NH3 (aq) + H2O (l)	——> <——	NH4^+(aq) + OH^-(aq)

• Nonelectrolytes: Nonelectrolytes are substances that dissolve in water, but do not generate any ions. Their solutions do not conduct electricity because of their total lack of ions. Common examples are:

o Ethanol (C₂H₅OH)

o Sucrose (C₁₂H₂₂O₁₁), also known as cane sugar.

4.3 The Composition of Solutions: Many important chemical reactions occur in solution. We need to be able to perform stoichiometric calculations on these reactions.

• Stoichiometry in Solutions: We need to know:
• The forms that the reactants and products take in solution, and
• The amounts of chemicals present in the solutions.

In the previous chapter, we learned to count the molecules in a reaction by weighing them, but when we deal with solutions, we need a different way of counting. The most convenient way to measure a liquid is to measure its volume. Thus, for example, if we had prepared a 1.000 liter solution containing 1.000 mol of sodium nitrate and then drew out a 50.00 mL portion of that solution, we would know that the 50.00 mL portion contains 0.05000 mol of sodium nitrate.

• Molarity: We can say that the above solution of sodium nitrate contains 1.000 moles of sodium nitrate in each 1.000 liters. Does this not look like a useful way to characterize a solution? So we define molarity as:

We can quantitatively describe our example solution as a 1.000 molar aqueous solution of sodium nitrate, and we can abbreviate that to 1.000 M NaNO₃.

• Sample Exercise 4.1 (p. 134): What is the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make up 1.5 L of solution? (The molar mass of NaOH is 40.00 g/mol.) We define some familiar symbols: m represents mass, MM represents molar mass, and n represents number of moles. And we need a new one: V for volume.

• Sample Exercise 4.2: (p. 134): Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl in enough water to make 26.8 mL of solution. (The molar mass of HCl is 36.46 g/mol.)
• Ion Concentrations: The above definition of molarity is based upon how the solution was prepared. What is the volume of the solution? How much solute did you weigh? It does not necessarily describe what is actually in the solution. For example, our 0.192 M solution of NaOH in Sample Exercise 4.1 does not contain NaOH molecules. The actual species and their molarities are:
• 0.192 M sodium ions.
• 0.192 M hydroxide ions.
• Suppose we had a 0.192 M solution of sodium sulfate (Na₂SO₄). The species in solution would be:
• 0.192 M sulfate ions.
• 0.384 M sodium ions (because each mole of sodium sulfate generates 2 moles of sodium ions).

	H2O
Na2SO4 (s)	——>	2Na^+(aq) + SO4^2-(aq)

• Sample Exercise 4.3 (pp. 134-5): What are the concentrations of each type of ion in the following solutions:
• 0.50 M Co(NO₃)₂
• 1 M Fe(ClO₄)₃
• Moles of Solute in a Given Volume: If we know the molarity of a solution, we can calculate how many moles of each of its components there are in a given volume. We need simply to determine the molarities of each component and multiply them by the volume. Remember that the units of the answer will be moles.
• Sample Exercise 4.4 (p. 135): How many moles of Cl^- ions are there in 1.75 L of 1.0 x 10^-3 M ZnCl₂?

	H2O
ZnCl2 (s)	——>	Zn^2+(aq) + 2Cl^-(aq)

• Standard Solutions: A standard solution is a solution whose concentration is accurately known. The procedure for preparing a standard solution is illustrated in Figure 4.10 form the text:

• Sample Exercise 4.6 (pp. 136-7): A chemist needs 1.00 L of an aqueous solution of K₂Cr₂O₇ (potassium dichromate) to analyze the alcohol contents of some samples of wine. How much K₂Cr₂O₇ must be weighed out to prepare the solution?
• The number of moles of K₂Cr₂O₇ required is:

• Then the mass of K₂Cr₂O₇ (molar mass = 294.20 g/mol) that must be weighed out is:

• Dilution: Instead of preparing large volumes of solutions with relatively low concentrations of solute, chemists will often instead prepare a smaller volume of a much more concentrated solution and dilute portions of it as needed. For example: Suppose the chemist in Sample Exercise 4.6 had needed 20 L of 0.200 M K₂Cr₂O₇. Rather than perform the steps in Figure 4.10 twenty times over, she might instead have prepared 1.00 L of a 4.00 M solution. Then each time she needed some 0.200 M solution for an analysis, she could accurately measure (by pipette) 50.00 mL of the 4.00 M solution into another 1 liter volumetric flask and add water to the 1.00 L mark. The key to dilution calculations is that:

Moles of solute before dilution = Moles of solute after dilution

In a dilution problem, you will know or be given information to calculate 3 of the 4 quantities in the line above.

• Example from page 137: We need 500. mL of 1.00 M acetic acid (HC₂H₃O₂) from a 17.4 M stock solution. What volume of stock solution do we need? (Work out on whiteboard.)
• Sample Exercise 4.7 (pp. 138-9): What volume 16 M sulfuric acid (H₂SO₄) must be used to prepare 1.5 L of a 0.10 M H₂SO₄ solution? (The solution in the text is a bit convoluted. We’ll work it out directly on the whiteboard.)

4.4 Types of Chemical Reactions in Solution

• Precipitation Reactions
• Acid-base Reactions
• Oxidation-reduction Reactions

These topics take up the rest of Chapter 4. We will cover them next week.