16.1 Properties of Acids and Bases

Let’s Review In the last chapter, we learned about solutions. We learned about why something dissolves. “Like dissolves like” Mass Percent and Molarity Dilutions Stoichiometry with solutions

In this chapter… We will still be learning about solutions. However, we are going to focus in on a particular type of solution. It all has to do with H + (a proton). Our main focus is to understand what happens in solutions when H + ions get involved.

We will learn about: The two models of acids and bases. The relationship of conjugate acid-base pairs. The strengths of acids. The ionization of water.

A brief history of acids and bases… Before 1850, acids were classified as substances with a sour taste; bases were classified by a bitter taste and slippery feel. In 1884, Svante Arrhenius became the first person to recognize the essential nature of acids and bases.

A brief history of acids and bases… Instead of classifying acids and bases on taste, he postulated that acids produced hydrogen ions (H + ) and bases produced hydroxide ions (OH - ).

A brief history of acids and bases… Although Arrhenius received a Noble prize based on his works in acids and bases, his model of bases was slightly off. It only allowed for the OH - ion. Soon after, Johannes Bronsted and Thomas Lowry (who worked independently of each other) reworked the models of acids and bases to the following:

A brief history of acids and bases… The Bronsted-Lowry model: An acid is a proton (H + ) donor, and a base is a proton acceptor. This is the model that we will use for the rest of the chapter.

Arrhenius vs. Bronsted-Lowry So far, we have learned about the two different models of acids and bases. AcidBase ArrheniusProduces H + Produces OH - Bronsted-LowryGives proton, gives H + Receives H +

More on the Bronsted-Lowry Model When an acid (“HA,” some substance that is ready to donate a H + ) dissolves in water, the reaction that occurs can be best represented by: HA + H 2 O  H 3 O + + A - Notice that after the reaction, the “H 3 O + ” (called the hydronium ion) molecule is now able to donate a H +. So now it is an acid. “A - ” is now able to accept a proton, so it is now a base.

More on the Bronsted-Lowry Model HA + H 2 O  H 3 O + + A - H 3 O + is called the conjugate acid - the substance formed when a proton is added to a base. A - is the conjugate base - the remaining substance when a proton is lost from an acid. One point of interest is how the acid and conjugate base relate to each other. In chemistry, “conjugate” means “differing by one proton.” A conjugate acid-base pair are related to each other by donating and accepting a single proton.

Skill: Finding Conjugates 1)Which of the following represent conjugate acid-base pairs? a) HF, F - b) NH 4 +, NH 3 c) HCl, H 2 O

Skill: Finding Conjugates 2)Write the conjugate base for each of the following: a) HClO 4 b) H 3 PO 4 c) CH 3 NH 3 +

Acid Strength Some acids are stronger than others. For instance, vinegar and citric acids are weak enough to use in foods. Whereas sulfuric acid is strong enough to be useful in car batteries. Question: What determines if an acid is strong or weak?

Acid Strength There are two ways to determine if an acid is strong or weak. Both ways have to do with the way the acid dissolves in water.

Acid Strength Remember this… HA + H 2 O  H 3 O + + A - Recall that H 3 O + is a conjugate acid and A - is a conjugate base. Since they are an acid and a base, they can react with one another! H 3 O + + A -  HA + H 2 O

Acid Strength So now the conjugate acid and conjugate base can re-form the parent acid and base. This reaction can occur “in both directions.” HA + H 2 O  H 3 O + + A - (We use a double arrow to represent both directions.)

Acid Strength HA + H 2 O  H 3 O + + A - This really represents a competition for the H + ions between H 2 O and A -.

Acid Strength If H 2 O has a higher attraction than A - (the conjugate base) for the H +, then we say that the “forward reaction predominates” and HA is completely ionized or completely dissociated. When an acid completely dissociates, we call it a strong acid.

Acid Strength HA + H 2 O  H 3 O + + A - If A - has a much larger attraction for H + than does H 2 O, most of the HA molecules remain as HA and therefore do not dissociate. This would be a situation where the “reverse reaction predominates,” which would mean that we have a weak acid.

Acid Strength

Strong Acids There are 4 common strong acids. Sulfuric acid, H 2 SO 4 (aq) Hydrochloric acid, HCl (aq) Nitric acid, HNO 3 (aq) Perchloric acid, HClO 4 (aq) The rest of the acids can be considered to be weak.

Types of Acids Diprotic - can furnish two protons (H 2 SO 4 ) Oxyacid – acidic proton is attached to an oxygen atom Organic acid – have a carbon atom backbone and commonly contain the carboxyl group

Water How can we tell if a solution is acidic, basic or neutral without knowing what is in it? The key to answering this question is found in what we can know about water.

Water Water also has the ability to act like an acid and a base (it is amphoteric). H 2 O + H 2 O  H 3 O + + OH - In pure water, only a tiny amount of H 3 O + and OH - exist. At 25 o C, the actual concentrations of ions are: [H 3 O + ] = [OH - ] = 1.0 X M

Water H 2 O + H 2 O  H 3 O + + OH - One of the most important things about water is that the mathematical product of the H 3 O + and OH - concentrations is always constant at 25 o C. K w = [H 3 O + ] [OH - ] = 1.0 X M K w is called the ion-product constant of water and it never changes, no matter what it contains. So, if [H + ] (shorthand for “[H 3 O + ]”) goes up, [OH - ] must go down.

Is it acidic or basic? If we add an acid to water, we have more H + ions, so [H + ] > [OH - ]. If we add a base to water, we have more OH - ions, so [H + ] < [OH - ]. If we have just pure water, [H + ] = [OH - ] This is called neutral.

Is it acidic or basic? In summary: Acidic solution, [H + ] > [OH - ] Basic solution, [OH - ] > [H + ] Neutral solution, [H + ] = [OH - ] In each case, K w = [H + ] [OH - ] = 1.0 X (25 o C).

Practice: Is it acidic or basic? Calculate [H + ] or [OH - ] at 25 o C and state whether the solution is neutral, acidic, or basic. a. 1.0 X M OH -

Practice: Is it acidic or basic? Calculate [H + ] or [OH - ] at 25 o C and state whether the solution is neutral, acidic, or basic. b. 1.0 X M OH -

Practice: Is it acidic or basic? Calculate [H + ] or [OH - ] at 25 o C and state whether the solution is neutral, acidic, or basic. c M H +

Practice: Is it acidic or basic? Is it possible for an aqueous solution at 25 o C to have [H + ] = M and [OH - ] = M?