We are starting with a recap of episode 35 and are defining acids and bases and their connection to pH and pOH.
We are starting with a recap of episode 35 and are defining acids and bases and their connection to pH and pOH (1:02). Taking a closer look at the pH scale, our episode discusses the use of logarithm for pH and pOH (1:53). This lays the foundation for the calculation of pH and pOH for strong acids and bases. For weak acids and bases the episode draws the connection back to equilibrium and introduces the acid/base ionization constant (3:10), which is needed together with the initial concentrations to calculate the pH of a weak acid/base (6:08).
Question: You have a 0.1 M sulfurous acid solution. If at equilibrium you have a [H+] = 8.4 x 10-3 M, what is the percent ionization?
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap pH and pOH.
Lets Zoom Out:
Unit 8 - Acids and Bases
Topic 8.2 and 8.3 - pH and pOH of strong and weak acids
Big idea - Structure and Properties
Stomach acid has a pH of 1, vinegar a pH of 3, coffee a pH of 5. That sounds confusing! Isn’t vinegar much more acidic than coffee? So why is the pH LOWER? Isn’t the pH the concentration of H+ ions? So many questions! Let’s take a closer look at the calculations of pH and pOH!
Let’s zoom in:
Let’s start with a recap of our recap episode 35 where we’ve already introduced pH and pOH: According to Brønsted-Lowry, acids are substances that donate protons - aka H+ ions. Bases are substances that accept H+ ions. To measure the concentration of hydronium ions, we calculate the pH, and to measure the concentration of hydroxide ions, we calculate the pOH. The pH and the pOH respectively are defined as the negative logarithm of the H+ and OH- ion concentration. Since the concentrations are usually less than 1 and often very small, scientists avoid using negative exponents by calculating the negative logarithm to get a rational number.
Let’s take a closer look at the pH scale, which we use far more often than the pOH scale. The pH scale is often displayed from 0 to 14, with 0 having a very high concentration of H+ ions and 14 having a very low concentration of H+ ions. Why is that? It is a function of using the negative logarithm. Let’s compare two different concentrations: a H+ concentration of 1 x 10-4 M and 1 x 10-3 M. It is easy to get confused: Which one is actually greater? 1 x 10-3 M is the same as 0.001M whereas 1 x 10-4 M equals 0.0001M. Therefore, the H+ ion concentration of 1 x 10-3 M is greater. The negative log of 1 x 10-3 M is 3, the negative log of 1 x 10-4 M is 4. Therefore, the lower the pH, the greater the concentration of H+ ions.
This already outlines some of our calculations: If we have the concentration of H+, we can take the negative log to calculate the pH. If we have the pH, we “reverse” our calculation by taking 10 to the -pH to get the concentration of H+.What if you have the molarity of a solution given, for example a 0.1 M solution of hydrochloric acid or a 0.2 M solution of ammonia? You will have to distinguish between strong and weak acids and bases for your calculation.
Strong acids completely dissociate in aqueous solutions. That means, if you dissolve 1 mole of hydrogen chloride to get a 1 liter solution with a concentration of 1 molar, the concentration of H+ as well as Cl- in the solution is 1 molar for each of them. So a 1 M solution of hydrochloric acid has a pH of 1. The strong acids are: hydrochloric, hydrobromic, hydroiodic, perchloric, sulfuric and nitric acid. All others are weak acids. That includes hydrofluoric acid! The same approach, but vice-versa applies to calculations for the pOH. The hydroxides of group I and II metal ions are strong bases, which completely dissociate. Therefore, if you dissolve 1 mol of sodium hydroxide in enough water to make a 1 L solution, your hydroxide concentration is 1 molar. This let’s us calculate the pOH to be 1. Know that we use the pH scale much more frequently, so what is the pH of a solution with a pOH of 1? As we’ve discussed in episode 35 when we introduced the Kw: pH and pOH together add up to 14. Therefore, with a pOH of 1, the pH is 13. This high pH again indicates that we have a very low concentration of H+ ions.
Now, what about all the acids and bases that are NOT strong? Weak acids and bases only dissociate to a small percentage. That means, the majority of species in the aqueous solution will be a molecule and not an ion. Therefore, the concentration H3O+ or respectively OH- will be very small. In these solutions, we have an aqueous equilibrium between the unionized acid and the H3O+ and conjugate base concentration. Let’s make an example: hydrofluoric acid is a weak acid. It’s dissociation equation is: HF ↔ H+ + F- . The Ka, acid ionization constant, for this reaction is 5.6 x 10-4 . As we know from Unit 7, that indicates that the reaction barely proceeds at all and supports our earlier statement that the majority of molecules are unionized. So how do we now determine the pH of a weak acid, if we cannot just use the molarity of the solution to determine the concentration of the H+ ions? Well, you already have the skills from Unit 7: We use the Ka as well as our initial concentrations and RICE tables to calculate the concentration of H+ at equilibrium. And then we can take the negative log to determine the pH. Tada!
The same approach also applies to weak bases - same same, but different: For weak bases we use the Kb, the base ionization constant. Setting up our equation showing the equilibrium between the weak base and the formation of OH- ions as well as the conjugate acid, we can use the initial concentration and RICE tables to determine the concentration of OH- ions at equilibrium. Taking the negative log will give us the pOH. Don’t forget the last step, when being asked about pH: pH is 14 minus pOH!
To recap:
To measure the concentration of hydronium ions, we calculate the pH and to measure the concentration of hydroxide ions, we calculate the pOH. Strong acids and bases completely dissociate. We therefore can use pH = -log [H+] and pOH = -log [OH-] as well as pH = 14-pOH to calculate the pH. Weak acids and bases have very small equilibrium constants, indicating that the majority of molecules are unionized. We therefore use the acid or base ionization constant and the initial concentrations to determine the concentrations of H+ and OH- at equilibrium and to calculate the pH.
Coming up next on the APsolute RecAP Chemistry Edition: acid-base titrations.
Today’s Question of the day is about percent ionization.
Question: You have a 0.1 M sulfurous acid solution. If at equilibrium you have a [H+] = 8.4 x 10-3 M, what is the percent ionization?