Acids and Bases

Chapter 11 Acids and Bases in our Lives

Acids and bases are important substance in health, industry, and the environment. One of the most common characteristics of acids is their sour taste. • Lemons and grapefruits taste sour because they contain acids such as citric and ascorbic acid (vitamin C). • Vinegar tastes sour because it contains acetic acid. Acids and Bases in our Lives

•We produce lactic acid in our muscles when we exercise. •Acid from bacteria turns milks sour in the products of yogurt and cottage cheese. •We have hydrochloric acid in our stomachs to help digest food and we take antacids, which are bases such as sodium bicarbonate, to neutralize the effects of too much stomach acid. Acids and Bases in our Lives

•In the environment, the acidity or pH of rain, water, and soil can have significant effects. •When rain becomes too acidic, it can dissolve marble statues and accelerate the corrosion of metals. •In lakes and ponds, the acidity of water can affect the ability of plants and fish to survive. •The acidity of soil around plants affect their growth. It can stop the plant from taking up nutrients through the roots Acids and Bases in our Lives

•The lungs and kidneys are the primary organs that regulate the pH of body fluids, including blood and urine. •Major changes in the pH of the body fluids can severely affect biological activities within the cells. Buffers are present to prevent large fluctuations. Chapter 11 – Acids and Bases

• 11.1 Acids and Bases • 11.2 Brønsted-Lowry Acids and Bases • 11.3 Strengths of Acids and Bases • 11.4 Dissociation Constants for Acids and Bases • 11.5 Dissociation of Water • 11.6 The pH Scale • 11.7 Reactions of Acids and Bases • 11.8 Acid-Base Titration • 11.9 Buffers 11.1 - Acids and Bases

Describe and name acids and bases. Acids

The term acid comes from the Latin word acidus which means “sour.” In 1887, the Swedish chemistry Svante Arrhenius was the first to describe acids as substances that produce hydrogen ions (H+) when they dissolve in water. Acids are Electrolytes

Because acids produce ions in water, they are also electrolytes (can conduct electricity). Hydrogen chloride dissociates in water to give hydrogen ions, H+, and chloride ions, Cl- :

It is the hydrogen ions that give acids a sour taste. Naming Acids

Acids have two common formats:

Binary acids: HnX Hn = some number of H’s x=nonmetals Examples: HCl, HBr, H, H2S…

Polyatomic acids: HnXOm XOm = polyatomic ion Examples: H2SO4, H3PO4, HClO4… Naming Acids

Binary acids: H X Change the ending of the nonmetal n to –ic and insert into the brackets. hydro[nonmetal –ic] acid hydro and acid do not change.

HCl

HBr

H2S Polyatomic Ion Review

2- More O’s = -ate SO4 2- Less O’s = -ite SO3

Chlorine can form 4 polyatomic ions with oxygen: - ClO4 - ClO3 - ClO2 ClO- Naming Acids

Polyatomic Acids: HnXOm [Polyatomic ion] acid -ate to –ic -ite to –ous

H2SO4

H3PO4

HClO3

Bases

• You may be familiar with some household bases such as antacids, drain cleaners, and oven cleaners. • According to the Arrhenius theory, bases are ionic compounds that dissociate into cations and hydrogen ions (OH-) when they dissolve in water. • They are electrolytes. Bases

Most Arrhenius bases are formed from a metal from Groups 1 or 2 and one or more hydroxides (OH-)

M(OH)n M=metal

(OH)n = 1 or more hydroxide group

Examples: LiOH, Ca(OH)2

The hydroxide ions give bases common characteristics such as a bitter taste or slippery feel. Naming Bases

Bases have the same names that we used for ionic compounds.

LiOH NaOH

Ca(OH)2

Al(OH)3 Chapter 11 – Acids and Bases

Identify the conjugate acid-base pairs for Brønsted- Lowry acids and bases. Arrhenius Acids and Bases

The definitions we gave in section 11.1 for acids and bases were first described by Arrhenius. So we call acids and bases described by H+ and OH- as Arrhenius acids and bases. Arrhenius acid: substances that produce H+ in water. Arrhenius base: substances that produce OH- in water. Brønsted-Lowry Acids and Bases

In 1923, a pair of scientists, J.N. Brønsted and T.M. Lowry expanded the definitions of acids and bases. The shortcoming of the Arrhenius definitions was that there were many molecules that didn’t have OH- groups that acted like bases. A new set of definitions describing Brønsted-Lowry acids and bases included a greater number of molecules. Brønsted-Lowry Acids and Bases

Brønsted-Lowry acid: a substance that donates a hydrogen ion, H+ Brønsted-Lowry base: a substance that accepts a hydrogen ion, H+

Arrhenius acid: produces H+ Arrhenius base: produces OH- + + H = H3O

• A free hydrogen, H+, does not actually exist in water. • Its attraction to polar water molecules is so strong that the H+ bonds to a + water molecules and forms a hydronium ion, H3O Brønsted-Lowry acid: donates H+ Brønsted-Lowry base: accepts H+ Brønsted-Lowry Acids

+ + - • HCl donates its H to water producing H3O and Cl

• By donating the H+, HCl is acting as the acid in this reaction. • By accepting the H+, water is acting as a base in this reaction. Brønsted-Lowry acid: donates H+ Brønsted-Lowry base: accepts H+ Brønsted-Lowry Bases

+ + - • Water gives an H to NH3 forming NH4 and OH

+ • NH3 acts as the base by accepting the H • Water acts as the acid by donating the H+ Water: a B-L acid and base

Water can act as both a Bronsted-Lowry acid or base depending on what it reacts with. Brønsted-Lowry acid: donates H+ Brønsted-Lowry base: accepts H+ Practice

Identify the reactant that is a Bronsted-Lowry acid and the reactant that is a Bronsted-Lowry base:

+ - HBr(aq) + H2O(l ) H3O (aq) + Br (aq) Brønsted-Lowry acid: donates H+ Brønsted-Lowry base: accepts H+ Practice

Identify the reactant that is a Bronsted-Lowry acid and the reactant that is a Bronsted-Lowry base:

- - CN (aq) + H2O(l ) HCN(aq) + OH (aq) Conjugate Acid-Base Pairs

According to Bronsted-Lowry theory, a conjugate acid-base pair consists of molecules or ions related by the loss of one H+ by an acid, and the gain of one H+ by a base. Every acid-base reaction contains two conjugate acid-base pairs because an H+ is transferred in both the forward and reverse directions. Conjugate Acid-Base Pairs

When an acid such as HF loses one H+, it becomes F-. HF is the acid, and F- is its conjugate base.

* The conjugate is always what is formed by donating or accepting H+. So it is always on the products side. Conjugate Acid-Base Pairs

+ + When the base H2O gains an H , its conjugate acid, H3O is formed. Conjugate Acid-Base Pairs

Now if we combine the two previous examples: Conjugate Acid-Base Pairs Amphoteric Substances

Water can act like an acid when it donates H+ or as a base when it receives H+ Substances that can act as both acids and bases are amphoteric. Water is the most common amphoteric substance and its behavior depends on the other reactant. Water will donate H+ when mixed with a base and will accept H+ when mixed with an acid. Amphoteric Substances

- Another example of an amphoteric substance is bicarbonate, HCO3 . - + - With a base, HCO3 acts as an acid and donates H to give CO3 . - + With an acid, HCO3 acts as a base and accepts H to give H2CO3 Practice

Identify the conjugate acid-base pairs in the following reaction:

- + HBr(aq) + NH3(aq) Br (aq) + NH4 (aq) Chapter 11 – Acids and Bases

Write equations for the dissociation of strong and weak acids; identify the direction of reaction. Strong vs Weak

In the process called dissociation, an acid or base separates into ions in water .

+ The strength of an acid is determined by the moles of H3O that are produced for each mole of acid that dissolves.

The strength of a base is determined by the moles of OH- that are produced for each mole of base that dissolves.

Strong acids and bases dissociate completely in water. Weak acids and bases dissociate only slightly, leaving most of the initial acid or base undissociated. Strong Acids

Strong acids are examples of strong electrolytes because they donate H+ so easily that their dissociate in water is essentially complete.

+ When HCl (a strong acid) dissociates in water, H is transferred to H2O. + - The resulting solution contains essentially only H3O and Cl .

• Thus one mole of a strong acid dissociates in water to yield one mole of H3O+ and one mole of its conjugate base. • We write the equation for a strong acid, such as HCl, with a single arrow. Weak Acids

Weak acids are weak electrolytes because they dissociate slightly in water, forming only a small + amount of H3O ions. When acetic acid dissociates in water, it donates the H+ to water. However, only part of the acetic acid molecules dissociate into ions. Most remain as molecules.

+ Thus one mole of a weak acid partially dissociates in water to give less than a mole of H3O - and C2H3O2 We write the equation for a weak acid in aqueous solutions with a double arrow to indicate that the forward and reverse reactions are at equilibrium. Strong and Weak Acids

There are only 6 common strong acids:

Hydroiodic acid HI Heavily regulated Hydrobromic acid HBr Used to make other molecules and extracting ore

Perchloric acid HClO4 Rocket fuel ingredient Hydrochloric acid HCl Stomach acid

Sulfuric acid H2SO4 Drain cleaner, lead-acid batteries

Nitric acid HNO3 Explosives ingredient Strong and Weak Acids

The rest are weak acids.

Diprotic Acids

Some weak acids, such as carbonic acid, are diprotic acids that have two H+, that dissociate one at a time.

For example, carbonated soft drinks are prepared by dissolving CO2 in water to form carbonic acid, H2CO3.

- + H2CO3 dissociates partially into HCO3 and H in water: + - H2CO3(aq) + H2O(l) H3O (aq) + HCO3 (aq)

- 2- + HCO3 is also a weak acid and will partially dissociate into CO3 and H - + 2- HCO3 (aq) + H2O(l) H3O (aq) + CO3 (aq) Diprotic Acids

Sulfuric acid, H2SO4, (a strong acid) is also a diprotic acid. + - H2OS4 will dissociate completely into H and HSO4 : + - H2SO4(aq) + H2O(l) H3O (aq) + HSO4 (aq)

- HSO4 is a weak acid and dissociates only partially: - + 2- HSO4 (aq) + H2O(l) H3O (aq) + SO4 (aq) Acid Summary

• A strong acid in water dissociates completely into ions. • A weak acid in water dissociates only slightly into a few ions but remains mostly as molecules.

- Strong acid: HI(aq) + H2O(l) H3O+(aq) + I (aq)

- Weak acid: HF(aq) + H2O(l) H3O+(aq) + F (aq) Bases

As strong electrolytes, strong bases dissociate completely in water. KOH(s) K+(aq) + OH-(aq)

Weak bases are weak electrolytes that are poor H+ acceptors and produce very few ions in solution.

+ - NH3(g) + H2O(l) NH4 (aq) + OH (aq) Bases in household products Direction of Reaction

There is a relationship between the components of each conjugate acid-base pair: Strong acids have weak conjugate bases. As the strength of the acid decreases, the strengths of the base increases. In any acid-base reaction, there are two acids and two bases. However one acid is stronger than the other acid. And one base is stronger than the other base.

+ - H3SO4(aq) + H2O(l) H3O (aq) + HSO4 (aq) Practice

By comparing their relative strengths, we can determine the direction of a reaction.

+ - H2SO4(aq) + H2O(l) H3O (aq) + HSO4 (aq) Practice

Which direction will the reaction favor?

2- - - CO3 (aq) + H2O(l) HCO3 (aq) + OH (aq) Practice

Which direction will the reaction favor?

+ - HF(aq) + H2O(l) H3O (aq) + F (aq) Chapter 11 – Acids and Bases

Write the expression for the dissociation constant of a weak acid or weak base. As we have seen, acids have different strengths depending on how much they dissociate in water.

Because the dissociation of strong acids in water is essentially complete, the reaction is not considered to be an equilibrium situation.

However, because weak acids in water dissociate only slightly, the ion products reach equilibrium with the undissociated weak acid molecules. Formic acid HCHO2, the acid found in bee and ant stings, is a weak acid. It + - dissociates in water to form hydronium ion, H3O , and formate ions CHO2 Writing Dissociation Constant Expressions

Because weak acids and bases reach an equilibrium when mixed in water, we can write an equilibrium constant expression (just like in ch. 10).

[Products] [D]d[C]c aA + bB cC + dD K = = a [Reactants] [A]a[B]b

Ka is called the acid dissociation constant. Practice

Write the equilibrium expression. + - HCHO2(aq) + H2O(l) H3O (aq) + CHO2 (aq)

* Only (aq) states are included in equilibrium expressions. (s) and (l) are ignored (including water). Writing Dissociation Constants

An equilibrium expression can also be written for weak bases: + - CH3-N2(aq) + H2O(l) CH3-NH3 (aq) + OH (aq)

[Products] Kb = = [Reactants]

* Only (aq) states are included in equilibrium expressions. (s) and (l) are ignored (including water). Dissociation Constants

• Just like in chapter 10, K’s less than 1 indicate that there is more reactant than product. • Which is in agreement of how we defined weak acids and weak bases. (Mostly molecules (reactants) and a small amount of ions (products)). • Strong acids and bases have very large K’s because its almost 100% dissociated. These K’s are not usually bothered with.

Chapter 11 – Acids and Bases

Use the water dissociation constant expressions to calculate the + - [H3O ] and [OH ] in an aqueous solution. In this section, we will use the dissociation constant expression and apply it to a very important equilibrium reaction: water reacting with itself. • In many acid-base reactions, water is amphoteric, which means tat it can act either as an acid or a base. • In pure water, there is a forward reaction between two water molecules that transfers H+ from one water molecule to the other. • One molecule acts as an acid by losing H+ and the other water molecule that gains H+ acts as the base. • Every time H+ is transferred between 2 water molecules, the products are Water one H3O+ and one OH-, which reacts in the reverse direction to re-form two water molecules. Water Dissociation Constant, Kw

+ - H2O(l) + H2O(l) H3O (aq) + OH (aq) Kw =

Experiments show that in pure water and 25°C, *ignore (s) and (l) [H3O+] = [OH-] =

If we plug the concentrations back into Kw: Kw = Neutral, Acidic, and Basic Solutions

The Kw applies to any aqueous solution at 25°C because all aqueous solutions + - contain H3O and OH . + - When [H3O ] and [OH ] in a solution are equal, the solution is neutral. However most solutions are not neutral; they have different concentrations of + - [H3O ] and [OH ]. Neutral, Acidic, and Basic Solutions

If acid is added to water, there is + an increase in [H3O ] and a decrease in [OH-], which makes it an acidic solution. If base is added to water, [OH-] + increases and [H3O ] decreases, which gives a basic solution. However for any aqueous solution, whether it is neutral, acidic, or basic, + - [H3O ][OH ] = 1.0 x 10-14

+ - Using Kw to calculate [H3O ] and [OH ]

+ - - If we know [H3O ], we can use Kw to calculate [OH ] or if we know [OH ] we + can use Kw to calculate [H3O ]. + - Kw = [H3O ][OH ] K K [OH-] = w [H O+] = w [H3O+] 3 [OH−] Practice

- -12 + A vinegar solution has a [OH ] = 5.0 x 10 M at 25°C. What is [H3O ] of the vinegar solution? Is the solution acidic, basic, or neutral? Practice

+ What is the [H3O ] of an ammonia cleaning solution with [OH-] = 4.0 x 10-4 M? Is the solution acidic, basic, or neutral? Chapter 11 – Acids and Bases