Chapter 20: Acids and Bases

Chapter 20: Acids and Bases

Describing Acids and Bases

History of theory for Acids and Bases

• Arrhenius, Svante – Swedish physical chemist (1859-1927) – one of the first

who attempted to scientifically describe acids and bases

• Arrhenius acid– substance that, when dissolved in water, produces

hydrogen ions (proton) • H+

• Arrhenius base– substance that, when dissolved in water, produces

hydroxide ions• OH-

Examples of Arrhenius Acid and Base

Arrhenius acid

Arrhenius base

Chapter 20: Acids and Bases-- Describing Acids and Bases --

Acids- Properties

• Aqueous solutions of acids are called electrolytes• Have a sour taste• Conduct electricity -- some well, some poorly• Cause some indicators, or chemical dyes, to change color• React with many metals to produce hydrogen gas• React with bases containing hydroxide ions to form salt and water• pH < 7

Bases- Properties

• Conduct electricity -- some well, some poorly• Taste bitter and feel slippery• Cause some indicators, or chemical dyes, to change color• React with acids containing hydroxide ions to form salt and water• pH >7


Names and Formulas of Acids• Acid takes the form of HX, where H is the hydrogen ion and X is a

monatomic or polyatomic ion

• Rules for naming– When the name of the anion (X) ends in “–ide”

- the acid name begins with hydro-- then, replace “-ide” with “ –ic” and add the word

“acid” to the end.

• Examples– HCl: hydrochloric acid– HBr: hydrobromic acid


Names and Formulas of Acids

– When the anion name ends in “-ate”• the acid is named by replacing “-ate” with “–ic” and adding the

word “acid” to the end.

• Examples– H2SO4: sulfuric acid (since it contains the sulfate ion)– HNO3: nitric acid (since it contains the nitrate ion)


Names and Formulas of Acids

– When the anion name ends in ”-ite”• the acid is named by replacing “-ite” with “-ous” , and

adding the word “acid” to the end.

• Examples:

– -H2SO3: sulfurous acid (since it contains the sulfite ion)

- HNO2: nitrous acid (since it contains the nitrite ion)


Names and Formulas of Bases

• Bases are named the same way as other ionic compounds – Example: NaOH= Sodium Hydroxide

- The cation is left the same

- The anion is found on our list of common polyatomic ions from the back of your periodic table

• ***One to know: NH3 = Ammonia

• Now, let’s practice!

Chapter 20: Acids and Bases

Hydrogen Ions and Acidity

Problems with the Arrhenius Theory

• Remember, Arrhenius said that….– acids were substances that dissociated into H+ ions in

solution (i.e. Hydrogen ion producer)– bases were substances that dissociated into OH- ions in

solution (i.e. Hyrdroxide ion producer)• In reality, this definition works OK for acids• This definition does not work for some bases

– baking soda NaHCO3 (a known base) does not produce OH- ions in solution.

Bronsted – Lowry Theory of Acids and Bases

• Acid – H+ donor (proton donor)

• Base – H+ acceptor (proton acceptor)

Water as Acid and Base• From the previous slide – water can act like acid (donate

proton) or base (accept proton),– this is called amphoteric

– In the presence of an acid, water will act like a base– In the presence of a base, water will act like an acid

• It is not surprising then that in pure water one molecule can donate a proton to another and water can act both like an acid and a base.

Chapter 20: Acids and Bases-- Hydrogen Ions and Acidity --

Hydrogen Ions from Water• Self-ionization

H2O + H2O H3O+ + OH-

Water Water Hydronium ion Hydroxide ion

– Hydrogen ions (hydronium ions) are often referred to as protons• These can be written as H+ or H3O+

– Self-ionization occurs very seldom, therefore the concentration of hydronium ions and hydroxide ions are low.

Concentrations of H+ and OH- in Pure Water

• Concentrations of hydrogen and hydroxide ions can be written several different ways

• For example, you have a sample of pure water whose concentration of H+ and that of OH- ions is 0.0000001 moles/ L, or 1.0 x 10-7 M – This can also be written as

• [H+] =1×10-7 M • [OH–]=1×10-7 M

– This means that these concentrations are equal in pure water.– Anytime the hydroxide and hydrogen ion concentrations are

equal, the solution is called a neutral solution


• For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14

[H+] x [OH-] = 1.0 x 10-14 M

• The product of these two ion concentrations is called the ion-product constant for water (Kw)


• Not all solutions are neutral. Solutions can be classified as acids or bases depending on their hydrogen and hydroxide ion concentrations

• Acids– [H+] > [OH-] – [H+] > 1.0 x 10-7 M

• Bases– [OH-] > [H+] – [H+] < 1.0 x 10-7 M


• Example:– If the [H+] in a solution is 1.0 x 10-5 M, is the

solution acidic, basic, or neutral? What is the [OH-] of this solution?


The pH Scale

• Expressing the concentration of hyrdogen and hydroxide ions in molarity can be difficult at times so scientists came up with another scale

• The pH scale– 0-14– Neutral solutions have a pH of 7– 0 is strongly acidic– 14 is strongly basic

• The pH scale is like the Richter scale, which measures earthquakes. The change of 1 unit is a tenfold

• Meaning- an earthquake with a tremor measuring 4.0 is 10x greater than one measuring 3.0.

• You can convert the concentration of hydrogen and hydroxide ions in molarity to the pH scale

pH = -log(H+)


• You must know all of these equations as well:

• Log (A * B) = logA + logB• Log (10x) = x• pH= -log[H+]• pOH= -log[OH-]• pH + pOH = 14


• Example– Find the pH of a solution with a hydrogen ion

concentration of 1.0 x 10-10M


• Example– The pH of a solution is 6.00. What is its

hydrogen-ion concentration?

The reality is…..

• pH is normally not a whole #• In this case we need a scientific calculator

• Example 1: What is the pH of a solution if [OH-] = 4.0 x 10-11M?– Solve by writing down your knowns and

unknowns from the problem

• Step # 1:– Known’s:

• [OH-]= 4.0 * 10-11M• KW= [H+] * [OH-] = 1.0 * 10-14

• pH = -log [H+]– Unknowns:

• pH=?• [H+]= ?

• I want to find pH. So, pH= -log [H+] but, we don’t know the [H+]. We must find that first:

• Step # 2: Find [H+]– [H+] = 1.0 * 10-14 = 0.25 * 10-3 M or 2.5 * 10-4

4.0 * 10-11

• Step # 3: Find pH– Use pH= -log [H+]

• -log (2.5 * 10-4)• Remember: log (A * B) = logA + logB• = - (log2.5 + log10-4)• = -(0.40) - (-4)• = -0.40 + 4• = 3.60 pH

• So, is this solution acidic or basic?

Now, you try:

• Example– What is the pH of a solution if [OH-] = 3.5 x 10-9 M?

One more:

• Calculate the pH of a solution with a [OH-] of 7.5 x 10-8 M

Let’s Consider this example:

• Fill in the table for a solution whose pH is 3.70.

pH pOH [H+] [OH-]

3.70

• Let’s do more practice!

pH pOH [H+] [pOH-]

3.70 10.3 1.0 * 10-3 1.0 * 10-10

Lets check your work

Chapter 20: Acids and Bases-- Types of Acids --

• Monoprotic acid– Acids that contain one ionizable hydrogenExample:

• HNO3

• Diprotic acid– Acids that contain two ionizable hydrogensExample:

• H2SO4

• Triprotic acid– Acids that contain three ionizable hydrogensExample:

• H3PO4

Chapter 21: Neutralization -- Neutralization Reactions --

• Acid-Base Reactions– An acid reacts with a base to produce water and salt

(generally)– Examples (strong acids reacting with strong bases)

• HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)• H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l)

• In the above examples there are equal number of hydrogen and hydroxide ions, creating a neutral solution.

• The final solutions have characteristics of neither an acidic or basic solution.


Neutralization reaction– Reaction in which an acid and a base react in an

aqueous solution to produce a salt and water

– A double-replacement reaction• You are expected to know how to predict the products if

given the reactants

– Neutralization will not necessarily occur between weak acids and/or weak bases (acids or bases that do not completely dissociate in solution).


• Titration– Process of adding a known amount of solution of a

known concentration to determine the concentration of another solution

– Titrations are helpful because they tell us when our solution is neutralized and no longer retains acidic or basic properties. (i.e tells us the exact concentration of base needed to neutralize an acid.)

Chapter 21: Neutralization -- Titrations--

Steps for completing a titration:1. A measured volume of an acid solution of unknown concentration is added

to a flask.2. Several drops of an indicator are added to the solution. 3. Measured volumes of a base of known concentration are mixed into the acid

until the indicator just barely changes color and maintains that color. This occurs at the “end point”.

4. When you add a base or acid of known concentration to this solution, add it incrementally until you reach the point at which the moles of H+ and OH– are equal. This is called the equivalence point.

– Examples for strong acids and strong bases– When titration is complete at the end point, the contents of the flask are

only salt and water. This is a neutral solution.

Titration Curve• Titration Curves

(This Curve is for Strong-Acid; Strong-Base Titration)

• The end point (equivalence point) has a pH at or very close to pH = 7.

Example Problem #1

• How many moles of sulfuric acid are required to neutralize 0.75 mol of potassium hydroxide?

Example Problem #1 – Answer Key

Titration Example #1

• The endpoint in a titration of 15.0 mL of 0.200M HCl with NaOH is reached when the volume of base added is 13.5 mL. Determine the molarity of NaOH.

Titration Example– Answer Key

Another Titration Example

Chapter 20: Acids and Bases

Documents

Transcript of Chapter 20: Acids and Bases