Acids and Bases / Acid Base Equilibrium

Unit 5 -Acids and Bases / Acid Base Equilibrium Text - Chapter 8

Review of Acids and Bases

Acids Bases

Naming Acids and Bases

Contains one or more hydrogen atoms

General formula

H – hydrogen atom n – number of hydrogen atoms (subscript) X – monoatomic or polyatomic anion

Examples Rules of Naming

1. When the name of the anion ends in “ide” (X), the acid is a binary acid, and the prefix is “hydro” and the ending is “ic”

2. When there is a polyatomic ion, that makes up (X), the acid is a ternary acid. When the polyatomic ion ends in “ate”, the ending for the acid is “ic”. If the ion ends in “ite”, the ending for the acid is “ous”

– If the name of the anion is “ate”, and the acid is “ic”, one less oxygen, the acid is “ous”, one more less oxygen, the acid is prefix “hypo” and ending is “ous” – If there is one more oxygen than the “ate” polyatomic ion, the name is, prefix “per” and ending “ic” – Some organic acids, you just have to memorize the name. Ex. Ethanoic Acid – CH3COOH or C2H4O2

Examples

Bases Named the same as ionic compounds Some you just have to memorize (ie. Ammonia – NH3)

Examples Acid-Base Systems – An Introduction (p. 528 – 532)

Arrhenius

The Arrhenius definition of acids and bases described an acid as a substance that dissociated in water to produce aqueous hydrogen ions, while a base was a substance that dissociated in water to produce aqueous hydroxide ions.

Ion – “wanderer”

+ - Example: HCl(aq)  H (aq) + Cl (aq)

+ - NaOH(s)  Na (aq) + OH (aq)

This approach works well for most substances that exhibit acid or base behaviour, but it does not work in all cases. For example, ammonia, is a base, but looks like an acid, due to its formula. We need another way of finding if it is an acid or a base.

Brønsted-Lowry

The Brønsted-Lowry definition focuses on the transfer of a proton (H+) so that an acid is defined as a proton donor, while a base is a proton acceptor.

+ - Example: HCl(aq) + H2O(l)  H3O (aq) + Cl (aq)

+ - NH3(aq) + H2O(l)  NH4 (aq) + OH (aq)

To act as an acid, a molecule or ion must have hydrogen in its structure that is involved in a polar covalent bond, while a base must have a lone pair of electrons in its valence shell.

Note that the role of water changes, acting as a base or an acid depending on the reaction. Any substance that can act as either an acid or a base is called amphoteric.

Examples Conjugate Acid-Base pairs

Let us consider the following reaction:

In the forward reaction NH3 is acting as a base since it is accepting a proton from H2O, while H2O is acting as an acid since it is donating a proton to NH3. In the reverse + - - reaction NH4 is acting as an acid since it is donating a proton to OH , while the OH is + acting as a base since it is accepting a proton from NH4 .

Any two species that differ only by a proton transfer are referred to as a conjugate acid- base pair. All Brønsted-Lowry acid-base reactions are an equilibrium between two conjugate acid-base pairs.

In each conjugate acid-base pair, the strengths of the acid and base are inversely proportional. So if the acid is a very good proton donor, the conjugate base is a very weak proton acceptor. If, however, the acid is a moderate proton donor, then the conjugate base will also have moderate strength.

Example: HCl(aq) is a very strong proton donor, readily giving off its proton to other - molecules. Conversely, Cl (aq) is a very weak proton acceptor, and will not accept protons from any other proton donors. As a result, when the equilibrium is established the system - will contain very few molecules of HCl(aq) and lots of Cl (aq) ions. H2CO3(aq) is only a - moderate proton donor, so its conjugate base, HCO3 (aq) also has moderate strength. As a result, measureable quantities of both species will be present at equilibrium.

The above is based upon acid / base strength, which we will go over next.

Homework – p. 532 #1-3 Acid / Base Strength (p. 534 – 540 and p. 551 – 562)

The degree to which an acid or base dissociates will determine the strength of the acid or base.

When a solution is made, there is a dynamic equilibrium between the bonded structure and the amount of ions that ionize. The greater the amount of ions in solution, the greater the degree of ionization and the stronger the acid or base. In other words, the Keq for a strong acid and base is greater than 1.

Based upon structure

- Greater the EN difference, the greater the ionization and dissociation, means the more “product” is formed, and more H+ or OH- goes into solution

- Therefore, stronger acid and base (Keq greater than 1)

- For polyprotic acids (ie. More than one hydrogen bonded to the structure), the first ionization is the strongest and the second and subsequent ionizations are weaker.

As we can see in the adjacent table, Hydrochloric acid, sulfuric acid and nitric acid are very strong. If we look at their structures and relative electronegativity differences, we see the hydrogen will be held loosely and is able to dissociate quite easily. The strong bases have greater electronegativity differences between the metallic ion and OH. Therefore, the hydroxide can separate and go into solution much more easily. For the others, the lower electronegativity difference, means that the ions will not be released into solution as readily, and are considered weak. + Keq for an acid is called Ka, and is the measure of how much of the acid ionizes (H or + H3O and X) and how much stays together (HnX)

+ - HX H + X

+ - H2XOy H + H XOy

- Keq for a base is called Kb, and is the measure of how much OH and + ion is in solution and how much stays together.

+ - M(OH)z M + Z( OH )

For Acids and Bases, the Keq greater than 1, means it dissociates 100%, and all of the reactant (acid or the base) goes to ions.

For example

3 H2SO4 Ka = 1.00 x 10

+ - H2SO4 H + H SO4

Therefore, if the concentration of the acid is 2.0 M, the concentrations of ions will be:

Therefore, when a strong acid or base dissociates, 100% turns into ions.

Example 2

What would be the OH- for a sodium hydroxide solution with a molarity of 0.5 M? What about weak acids and bases?

We need to solve for the concentrations of ions, using Ka or Kb values because they do not dissociate 100%

Ka or Kb values are less than 1, favouring the acid or base to stay together and little ionize. They may only dissociate 50% or 10%, leaving the majority of the acid or base as a “whole” and very little in ion form. Example

What is the [H+] concentration in a 1.0 M solution of carbonic acid? (Ka = 4.3 x 10-7)

Why are we finding the H + or OH - ?

Homework

1. You have a 0.3 M solution of hydrobromic acid? What is the hydronium ion concentration?

2. What is the [OH-] for a 1.0 M solution of ammonia? (Kb = 5.56 x 10 -10)

3. What are the concentrations of the ions produced, in a 0.5 M solution of acetic acid? (Kb = 1.80 x 10 -5) Relative Acidity/Alkalinity of Solution - pH

The scale used to compare the acidic or basic nature of solutions is the pH scale which was developed by Søren Sørensen in 1909.

+ + -pH pH = -log[H3O ] or [H3O ]=10

Since water is amphiprotic, i.e. it can act as either a proton donor or a proton acceptor, it is possible for water to react with itself. In a sample of pure water, a small number of water molecules with undergo a process of self-ionization by the following equation.

+ - 2 H2O(l) ⇄ H3O (aq) + OH (aq)

+ - [H3O ][ OH ] K =  2 [H2O]

Again, since [H2O] is fairly constant, it can be eliminated from the equations to leave

+ - -14 Kw = [H3O ][ OH ] = 1.0 x 10 , the dissociation constant for water

+ - -7 In a neutral solution, [H3O ] = [ OH ] = 1.0 x 10 , pH = 7

+ - + -7 In acid [H3O ]  [ OH ], [H3O ] 1.0 x 10 , pH < 7

+ - + -7 In base [H3O ] < [ OH ], [H3O ] < 1.0 x 10 , pH  7

The general range of pH values runs from 0 (very acidic) to 14 (very basic) Examples

What is the pH of a 0.1 M solution of hydroiodic acid? What is the hydroxide concentration?

If a sodium hydroxide solution has an [OH-] of 4.50 x 10-9 M, what is the pH of the solution?

What is the pOH of a solution with a pH = 4.56? What is the hydronium ion concentration?

Homework p. 540 #9 and 10 p. 546 #12 and 13 p. 549 #17, 18. and 19

-4 1. What is the pH of a 0.1 M solution of hydrofluoric acid? (Ka = 3.50 x 10 ) What is the percent ionization?

2. A solution of propanoic acid has a pH of 2.78. Calculate the percent ionization of a 0.05 M solution.

-10 3. Calculate the pH of a 0.1 M solution of ammonia. (Kb = 5.56 x 10 )

Acid-Base Indicators (p.608 – 610)

An acid-base indicator is a weak acid or a weak base that changes color as it accepts or donates a proton. The protonated form of the indicator is a different color than the unprotonated form of the indicator. An indicator does not change color from pure acid to pure alkaline at specific hydrogen ion concentration, but rather, color change gradually occurs over a range of hydrogen ion concentrations. This range is termed the color change interval. It is expressed as a pH range.

A commonly used acid-base indicator that you may be familiar with is litmus. It is often absorbed onto filter paper. The resulting piece of paper or solution with water becomes a pH indicator (one of the oldest), used to test materials for acidity. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic (i.e. alkaline) conditions, the color change occurring over the pH range 4.5-8.3 (at 25°C). Neutral litmus paper is purple in color

+ - HLit + H2O ⇄ H3O + Lit Acid form Base form Red Blue

At any pH below 4.5 we will see only red, while at any pH above 8.3 we will see only blue. In the pH range between 4.5 and 8.3, the red and blue forms of the indicator will be in approximately equal amounts and we will see the intermediate color purple.

Movie – Acid – Base Indicators Neutralization and Titration (p. 595 – 608)

If we have high levels of acid in our stomach, we take an antacid (base) to control it. In our small intestine, the acidic chyme, is neutralized by the bile (basic) to make sure we do not ulcerate our intestine

Why does it neutralize? We make a salt and water, that does not necessarily work out to a pH of 7 (more later) or we could bring the solution to a pH of 7.

If you react a strong acid and a strong base, the ions in solution will cancel each other out, producing a neutral solution. The products are always, regardless of the product pH, a salt and water.

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

The point where the number of moles of hydronium ion equals the number of moles of hydroxide ion, is called the equivalence point. (not necessarily pH =7)

Salt - The compound formed by the cation of the base bonding with the anion of the acid.

Titration

The process of adding a known amount of solution of known concentration to determine the concentration of another solution.

Titration is a general class of experiment where a known property of one solution is used to infer an unknown property of another solution. In acid-base chemistry, we often use titration to determine the pH of a certain solution or determine the concentration of a solution. The equivalence point of the titration is the point at which we have mixed equal numbers of moles of acid and base. For a strong acid-strong base titration, the equivalence point will occur at a pH of 7, and a salt and water will be the occurring products. However if the relative strengths of the acid and base are different, the equivalence point may not occur when the solution is neutral.

For a weak acid-strong base titration, the equivalence point will occur at a pH greater than 7, while a strong acid-weak base titration will have an equivalence point at a pH less than 7.

An indicator is used when titration is done. When the color changes, this is called the end point, which is the point of neutralization. Another way would be to use a pH meter, to measure the overall pH of the solution.

A setup for the titration of an acid with a base is shown to the right. We use this instrumentation to calculate the amount of unknown acid in the receiving flask by measuring the amount of base, or titrant, it takes to neutralize the acid. Titration Curves

A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the y-axis. The titration curve serves to profile the unknown solution. In the shape of the curve lies much chemistry and an interesting summary of what we have learned so far about acids and bases.

The titration of a strong acid with a strong base produces the following titration curve:

Note the sharp transition region near the equivalence point on the curve. Also remember that the equivalence point for a strong acid-strong base titration curve is exactly 7 because the salt produced does not undergo any hydrolysis reactions.

However, if a strong base is used to titrate a weak acid, the pH at the equivalence point will not be 7. There is a lag in reaching the equivalence point, as some of the weak acid is converted to its conjugate base. A solution composed of a mixture containing a weak acid and its conjugate base is known as a buffer. Buffers are solutions that resist changes in pH when small amounts of acid or base are added. In the diagram below we see the resultant lag that precedes the equivalence point, called the buffering region. In the buffering region, it takes a large amount of NaOH to produce a small change in the pH of the receiving solution.

The titration curve shown above is for a diprotic acid such as H2SO4 and is not unlike two stacked . For a diprotic acid, there are two buffering regions and two equivalence points. This proves the earlier assertion that polyprotic acids lose their protons in a stepwise manner.

Homework p. 599 #1 and 2 Salts

 Salts are the combination of the cation from the base and the anion from the acid and are the products of neutralization reactions

 Salts can be acidic, basic or neutral, depending on the strength of the acid and base that formed it.

 Buffers are an equilibrium condition which consists of the weak acid and conjugate base (salt) in solution, or a weak base and its conjugate acid (salt) keeping pH stable

 Generally, if a strong acid reacts with a strong base, the resulting salt will be neutral (pH=7) (ie. The equivalence point is 7)

 For salts formed from weak acids with a strong base, or weak bases with a strong acid, the salt will not be neutral

 This is called by salt hydrolysis, as the cations or anions from a dissociated salt remove or add hydrogen ions to water, creating either H+, or OH- in solution

 Another way to determine the acidity or basicity of a salt is to look at the net ionic equation and remembering that strong bases and acids dissociate 100%, while weak acids and bases do not.

 In general  Acidic salts produce positive ions that release protons into water  Basic solutions produce negative ions that attract protons from water

Examples of the Hydrolysis of Salts Buffers (p. 615 – 620)

Many aqueous solutions resist a change in pH upon addition of small amounts of acid or base. Such solutions, called buffer solutions, are said to be buffered.

Human blood, for example, is a complex aqueous medium with a pH buffered at about 7.4. Any significant variation of the pH from this value results in a severe pathological response and, eventually, death. As another example, the chemical behaviour of seawater is determined in very important respects by its pH, buffered at about 8.1 to 8.3 near the surface. Addition of a small amount of an acid or base to either blood or seawater does not result in a large change in pH.

A buffer contains the weak acid and one of its salts (anion – conjugate base) or the weak base and one of its salts (cation – conjugate acid). Pure water is not a buffer, as when you add acid (H+) or base (OH-), the concentrations increase, changing the pH.

A buffer is like a sponge

When hydrogen ions are added, they are absorbed by the negative ion, forming the “whole” weak acid, that does not dissociate 100%, lowering the acidity, and raising the pH to neutral

When hydroxide is added, they react with the acid to form the negative ion and water, lowering the basicity, and lowering the pH to neutral

To understand how a buffer works, let’s consider a solution of HC2H3O2 and NaC2H3O2. Ionization of HC2H3O2 is governed by the following equilibrium reaction:

+ - HC2H3O2(aq) ⇄ H (aq) + C2H3O2 (aq)

- The C2H3O2 in this equilibrium comes from both HC2H3O2 and NaC2H3O2. This mixture can either react with surplus H+ ions or release them, according to circumstances. For example, if a small quantity of acid is added to the solution, the equilibrium shifts to the left; acetate ion reacts with the added H+. The solution thereby limits pH change due to added acid. On the other hand, if a small quantity of a base is added, it reacts with H+. This reaction causes the equilibrium to shift to the right; HC2H3O2 dissociates to form more H+. The solution thereby also resists change in pH due to added base.