1 Chapter 19 Acids, Bases, and Salts

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1 1 Chapter 19 Acids, Bases, and Salts ACID-BASE THEORIES Acids and bases are all around us and part of our everyday life (ex. bodily functions, vinegar, carbonated drinks, citrus fruits, car batteries, antacids, cleaning products, etc.). Acids and bases have the following properties: Acids Tart or sour taste Electrolytes when in aqueous solutions (some are strong electrolytes, some are weak) Cause blue litmus paper (an indicator) to turn red Produces water and a salt when reacting with a base containing hydroxide Produces H2(g) when reacting with many metals Bases Bitter taste Electrolytes when in aqueous solutions (some are strong electrolytes, some are weak) Cause red litmus paper to turn blue Produces water and a salt when a base that contains hydroxide reacts with an acid Feel slippery (1) ARRHENIUS ACIDS AND BASES (1887) Svante Arrhenius proposed that acids are hydrogencontaining compounds that ionize to yield hydrogen ions (H + ) in aqueous solution and that bases ionize to yield hydroxide ions (OH - ) in aqueous solution. Acids that contain one ionizable hydrogen (ex. HNO3(aq), HCl(aq)) are called monoprotic acids. If an acid contains two and three ionizable hydrogens (ex. H2SO4(aq), H3PO4(aq)) they are known as diprotic and triprotic acids, respectively. Not all compounds that contain hydrogen are acids (ex. CH4(g)). Also, not all hydrogens of an acid may be released as hydrogen ions. Only hydrogens in very polar bonds are ionizable (ie. when they are bonded to a highly electronegative atom). Dissolving in water will release hydrogen because it is stabilized by its attraction to water. Examine the example below on the left and notice the hydronium ion that is created. Then examine the images on the right and consider why methane is different than acetic acid (ethanoic acid). Can you see why ethanoic acid would be considered a monoprotic acid? As stated earlier, Arrhenius bases release hydroxide ions when dissolved in water. Some bases are more soluble in water than others (see the solubility table on the back of the periodic table). When Group 1A elements (alkali metals) are exposed to water they react violently and produce basic solutions and hydrogen gas (as seen below). The Arrhenius definition of an acid and base unfortunately does not include certain substances that have acidic or basic properties (ex. Na2CO3(aq) and NH3(aq) display basic properties). (2) BRONSTED-LOWRY ACIDS AND BASES (1923) Johannes Bronsted and Thomas Lowry independently proposed that an acid is a hydrogen-ion donor and that a base is a hydrogen-ion acceptor.

2 The example below shows two substances that would not be considered a base and an acid according to the Arrhenius definition. Examine why these would be considered a Bronsted-Lowry base and acid. Hydrogen ions are transferred from water to ammonia and causes the hydroxide-ion concentration to be greater than it is in pure water. As a result, solutions of ammonia are basic. 2 Gases become less soluble in water as temperature rises (consider collision theory and why this would be the case!). Consequently, increasing the temperature of an aqueous solution of ammonia releases ammonia gas. According to Le Chatelier s principle, the equilibrium position would shift to the left because ammonium will give up a hydrogen and donate it to the hydroxide-ion. That means that during this shift in equilibrium, ammonium is a Bronsted-Lowry acid and hydroxide is a Bronsted-Lowry base. A conjugate acid is the particle formed when a base gains a hydrogen ion. A conjugate base is the particle that remains when an acid has donated a hydrogen ion. The Bronsted-Lowry theory also applies to acids (as seen below in the dissociation of hydrogen chloride in water). Notice that water gains the hydrogen ion and becomes the hydronium-ion. Sometimes water accepts a hydrogen ion and other times it donates it. A substance that can act as both an acid and a base, like water, is said to be amphoteric. (3) LEWIS ACIDS AND BASES (~1930s) Gilbert Lewis proposed that an acid accepts a pair of electrons during a reaction while a base donates a pair of electrons. This is the most general theory of the three. A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. A Lewis base is a substance that can donate a pair of electrons to forma a covalent bond. A hydrogen ion (Bronsted-Lowry acid) can accept a pair of electrons and, therefore, it is also a Lewis acid. A Bronsted- Lowry base, or a substance that accepts a hydrogen ion, must have a pair of electrons available and, therefore, is also a Lewis base. Keep in mind, however, that the Lewis definition also includes some compounds not classified as Bronsted-Lowry acids or bases. Summary: Conceptual Problem 19.1 p. 593, Practice Problems 1 and 2.

3 HYDROGEN IONS AND ACIDITY Water molecules are highly polar and are moving constantly. Sometimes they can collide with one another and transfer a hydrogen ion to produce OH - and H3O + (the hydroxide ion and hydronium ion, respectively). We call this the self-ionization of water and it can be written as a simple dissociation equation: H2O(l) H + (aq) + OH - (aq). In water, H + is always joined to water molecules to make H3O +. Hydrogen ions in aqueous solutions are given several names that are all acceptable to use: protons, hydrogen ions, or hydronium ions. 3 The self-ionization of water is minimal at 25 C. In pure water at this temperature, [H + ] and [OH - ] are only 1 x 10-7 M. Any solution that involves and equal concentration of [H + ] and [OH - ] is classified as a neutral solution. Given this information, we have established the ion-product constant for water (K w) as: Not all solutions are neutral. When considering Le Chatelier s Principle and the net ionic equation for aqueous solutions (on the right), dissolving a common ion will change the concentration of the ion and will drive the equilibrium position toward the production of more water. Another way to look at this, for example, is if you increase the [H + ] then you decrease the [OH - ] in order to produce the water. The alternative is also true if you increase [OH - ] then the [H + ] will decrease in order to produce more water. When H 2O HCl(g) H + (aq) + Cl - (aq), the [H + ] becomes greater than [OH - ] and this defines an acidic solution. This means that [H + ] is greater than 1 x 10-7 M. When H 2O NaOH(s) Na + + OH - (aq), the [OH - ] becomes greater than [H + ] and this defines a basic (or alkaline) solution. This means that the [H + ] is less than 1 x 10-7 M. Sample Problem 19.1 p.596, Practice Problems 9 and 10. It is easier to express [H + ] using the ph scale. The ph of a solution is the negative logarithm of the [H + ] and is represented by. The ph scale runs from 0 to 14 and represents whether or not a solution is acidic (<7), basic (>7), or neutral (7). In a neutral solution, the [H + ] is 1 x 10-7 M and we say that the ph is 7. Analyze the information here on the right. A solution in which the [H + ] is greater than 1 x 10-7 M ends up having a ph less than 7 and is considered acidic. The ph of pure water or a neutral aqueous solution is 7.0. A solution with a ph greater than 7 is basic and has a [H + ] that is less than 1 x 10-7 M.

4 4 The poh of a solution equals the negative logarithm of the [OH - ]: There is a simple relationship between ph and poh: For ph calculations, the [H + ] should be expressed in scientific notation. As for significant figures, consider the follow example: If [H + ] = M, then [H + ] = 1.0 x 10-3 M and since it has two sig figs, the ph would take them into account to dictate how many digits to place after the decimal (ie. ph = 3.00). Sample problem 19.2 p. 599, Practice Problems 11 and 12. Sample problem 19.3 p. 600, Practice Problems 13 and 14. Sample problem 19.4 p. 601, Practice Problems 15 and 16. An indicator (HIn) is an acid or a base that undergoes dissociation in a known ph range; it is a valuable tool for measuring ph because its acid form and base form have different colors in solution. The following generalized equation represents the dissociation of an indicator: In this example, the acid form dominates the dissociation at a low ph (high [H + ]) and the base form dominates the equilibrium at high ph [OH - ]. For each indicator at 25 C, the change from dominating acid form to dominating base form occurs in a narrow range of approximately two ph units; although the color of the solution is a mixture of the colors of the acid and the base form within this narrow range, it can give you a rough estimate of the ph. You could get a more precise estimate of the ph of the solution by repeating the experiment with indicators that have a different ph ranges for their color changes. The effectiveness of indicators and be reduced if experimentation is not done at 25 C, if the solution being tested isn t colorless, and if there are dissolved salts present; because of this, ph indicator strips and ph meters (even better!) are also available and can be more effective.

5 STRENGTHS OF ACIDS AND BASES In general, a strong acid will completely ionize in aqueous solution and leave a high concentration of H3O +. Weak acids will only slightly ionize, meaning only a small percentage of the acid is ionized at any instant and thus there is a low concentration of H3O +. 5 Example of a strong acid in water: Example of a weak acid in water: An equilibrium constant expression can be written for an acid. Consider ethanoic acid: For dilute solutions, the concentration of water is a constant and can therefore be combined with Keq to give an acid dissociation (or ionization) constant (K a). Weak acids have a small Ka. The stronger an acid is, the larger its Ka value. Based on the Ka expression, can you rationalize why this would be the case? Diprotic and triprotic acids lose their hydrogens one at a time and so each ionization reaction has a separate Ka. Just as there are strong and weak acids, there are also strong and weak bases. Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution. Weak bases react with water to form the hydroxide ion and the conjugate acid of the base. Consider ammonia a weak base: Notice the similarities to the acid dissociation constant. The base dissociation constant is small for weak bases and large for strong bases. Again, given the expression, why does this trend make sense? What is the difference when saying that an acid or base is concentrated/dilute compared to saying they are strong/weak? The words concentrated and dilute are considering the number of moles of the substance in one liter of water. The words strong and weak are referring to the extent of ionization/dissociation of the acid/base. Concentrated/dilute and strong/weak cannot be used interchangeably. Calculating dissociation constants: Sample Problem 19.5 p.610, Practice Problems 22 and 23.

6 NEUTRALIZATION REACTIONS In general, reactions in which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions. The substances must be mixed in the mole ratios specified by the balanced chemical equation. A neutralization reaction is also a way to prepare pure samples of salts. Recall that salts are compounds consisting of an anion and a cation. 6 When an acid and a base are mixed, the equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions. Depending on the balanced chemical equation, sometimes an acid and base react in a 1:1 ration and other times they don t. Examine the samples of neutralization reactions below: Sample Problem 19.6 p. 614, Practice Problems 30 and 31 By using the appropriate acid-base indicator, you can determine the concentration of acid (or base) in a solution by performing a neutralization reaction. Phenolphthalein is often preferred for this sort of investigation because it goes from colorless in an acidic solution to a faint pink as soon as the solution is slightly basic. The process of adding a known amount of solution of known concentration to determine the concentration of another solution is called titration. The following steps are involved in a titration: (1) A measured volume of an acid solution of unknown concentration is added to a flask. (2) Several drops of the indicator are added to the solution while the flask is gently swirled. (3) Measured volumes of a base of known concentration (called the standard solution) are mixed into the acid using a burette until the indicator just barely changes color. Titration continues until the point that the indicator changes color to show that neutralization has just occurred (we say that the end point has been reached). Just as the concentration of an unknown acid can be determined using an indicator and the titration of a standard base, the concentration of an unknown base can be determined using a standard acid in the same way. The ph (or titration) curve on the right shows how the ph of a solution changes during the titration of a strong acid (HCl) with a strong base (NaOH). Notice how the ph of the initial solution is low. As the base is added, the ph increases because some of the acid is neutralized. As the titration approaches the end point, the ph increased dramatically as H + are used up. At the equivalence point, the beaker will consist of only H2O, the salt (NaCl), and a trace of indicator. Once past the equivalence point (so at the end point and beyond), additional base produces a further increase of ph. Sample Problem 19.7 p. 616, Practice Problems 32 and 33. SALTS IN SOLUTION In the context of neutralization reactions, a salt consists of an anion from an acid and a cation from a base. Most salt solutions are neutral, but some can be acidic and others can be basic. For example. A solution of ammonium chloride is acidic and a solution of sodium acetate is basic (consider the Bronsted- Lowry definition of acids and bases). Consider the titration of ethanoic acid (a weak acid) with sodium hydroxide (a strong base) as the titrant: Notice that the ph of the equivalence point is 8.7 (basic) for the neutralization of a weak acid with a strong base, whereas the ph of the equivalence point for a strong as and strong base is 7 (neutral). Why is there a difference?

7 For a strong acid-base titration, the ph at the equivalence point is 7. For the titration of a strong acid with a weak base (or vice versa), the solution can end up being slightly acidic or slightly basic. This difference exists because the cations or anions of certain dissociated salts remove hydrogen ions from or donate ions to water in a process called salt hydrolysis. Hydrolyzing salts are usually derived from strong/weak acid-base titration. In general, salts that produce acidic solutions contain positive ions that release protons to water to make hydronium ions. Salts that produce basic solutions contain negative ions that attract protons from water and render hydroxide ions. Salts that come from strong acids and strong bases are not found to be hydrolyzing salts. 7 Ex. 1: Weak acid titrated with a strong base (CH 3COOH (aq) and NaOH (aq)) Ex. 2: Weak base titrated with a strong acid (NH 3(aq) and HCl aq)) The following rules may be a helpful reminder of the theory discussed. Make sure you are clear about WHY these rules are applicable! A buffer is a solution in which the ph remains the relatively constant when small amounts of acid or base are added. A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts. They react with any hydroxide ions or hydrogen ions added to the solution. For example, when an acid is added to the solution represented on the right, the ethanoate ions accept hydrogen from the acid to produce ethanoic acid. Because this weak acid doesn t ionize extensively in water, the ph doesn t change much. If a base was added to solution, the ethanoic acid donates a hydrogen to hydroxide to produce water and the ethanoate ion. Since the ethanoate ion isn't a strong enough base to accept hydrogen ions extensively, the ph does not change very much. If too much acid or base are added then the buffer capacity of a solution has been met, and it will no longer be able to control the ph. The buffer depends on the number of moles of the weak acid/ base and if too much acid or base is added, there will no longer be any left to react with the excess hydrogen-ion or hydroxide-ion. Conceptual Problem 19.2 p.622, Practice Problems 38 and 39.