ACIDS AND BASES



1. SC7. Students will characterize the properties that describe solutions and the nature of acids and bases.

1. b. Compare, contrast, and evaluate the nature of acids and bases:

2. • Arrhenius, Bronsted-Lowry Acid/Bases

3. • Strong vs. weak acids/bases in terms of percent dissociation

4. • Hydronium ion concentration

5. • pH

ACIDS AND BASES

Properties

Taste

Acids taste sour or tart.

Bases (water solutions) taste bitter.

B. Touch

Dilute acids feel like water. Stronger acids are corrosive.

Mild bases feel smooth and slippery. Strong bases are corrosive.

Reaction with metals

Acids react vigorously with metals, sometimes producing a gas.

Mg + HCl ( MgCl2 + H2 (g)

Reaction with Carbonates/Hydrogen Carbonates

Acids react vigorously with carbonates or hydrogen carbonates, producing a gas.

NaHCO3 + HC2H3O2 ( Na C2H3O2 + H2O + CO2

Electrical conductivity

Pure water is an extremely poor conductor of electricity.

An aqueous solution of a strong acid or base conducts electricity well.

Acids and bases are examples of electrolytes. (Electrolytes ionize when they dissolve in water. A solution containing electrolytes conducts electricity.)

Indicators

An acid-base indicator is a substance that turns one color in an acidic solution and another color in a basic solution.

Ex. Litmus (acid is red and base is blue), phenolphthalein, methyl red

Neutralization

If an acid and a base are mixed together, a two new compounds are formed. The properties of the acid and base are destroyed; the acid and base are neutralized. One of the products formed is always an ionic compound and is called a salt.

DEFINING ACIDS AND BASES

To truly understand acids and bases, you must study them at the molecular level. Their properties come from their structure and composition of molecules.

Acids basically are substances that have more H+ ions than OH- ions. Bases have the opposite and when something is neutral, it has equal amounts of H+ and OH-. Usually, acids and bases are dissolved in water; water provides both H+ and OH- ions. Sometimes you will see H3O+ (hydronium ion) instead of H+ (hydrogen ion).

The Arrhenius Definition:

Acids and bases can be understood in terms of the ions they release when they dissolve in water.

1. An acid is a substance that dissociates in water to produce hydrogen ions (H+).

2. A base is a substance that dissociates in water to produce hydroxide ions (OH-).

HCl + NaOH ( NaCl + H2O

H+ + Cl- + Na+ + OH- ( Na+ + Cl- + H2O

Arrhenius acids include HCl, HNO3, H2SO4, HC2H3O2, H2CO3, and H3PO4.

Arrhenius bases include NaOH, KOH, Mg(OH)2, Ca(OH)2, and Ba(OH)2. Notice they are all hydroxides.

With Arrhenius acids and bases, an acid-base neutralization reaction always produces water and a salt. Metals and acids produce a salt and hydrogen gas.

The Brønsted-Lowry Definition:

A redefined definition of acids and bases that is independent of how they act in water and it focuses solely on H+ ions. Notice that OH- ions are not even addressed. NH3, a base, was not even addressed by Arrhenius’ definition.

1. An acid is any substance that can donate H+ ions.

2. A base is any substance that can accept H+ ions.

HCl (g) + H2O (l) ( H3O+ (aq) + Cl - (aq)

The hydronium ion is usually formed from H+ and H2O

Conjugate Acids and Bases

When an acid loses an H+ ion it becomes a conjugate base. For example, hydrochloric acid (HCl) loses an H+ ion to become its conjugate base Cl-1. The conjugate base of water, H2O, is the hydroxide ion, OH-1. Likewise, when a base gains an H+ ion, it becomes its conjugate acid. For example, ammonia, NH3 gains an H+ to become its conjugate acid NH4+. And the conjugate acid of OH-1 is H2O. Generally speaking, the stronger the acid, the weaker the conjugate base and vice versa.

VIDEODISC: Defining Acids and Bases

1. What is the difference between Arrhenius base and a Brønsted-Lowry base?

2. What is the conjugate base of HCl? Is it strong or weak?

3. Name three acids that are 100% ionized in water?

Amphoteric – a substance that acts as both an acid and a base.

HF + H2O ( H3O+ + F- NH3 + H2O ( NH4+ + OH-

base acid

MONOPROTIC AND POLYPROTIC ACIDS

Monoprotic acids have only one hydrogen to lose. HF, HCl, HClO4

Polyprotic acids have two or more hydrogen to lose. H2SO4, H3BO3

Polyprotic acids ionize in steps, only one hydrogen at a time. Each time the acid gets weaker.

H3BO3 + H2O ( H2BO3- + H3O+

H2BO3- + H2O ( HBO3-2 + H3O+

HBO3-2 + H2O ( BO3-3 + H3O+

ANHYDRIDES

Some oxides become acids or bases by adding the elements contained in water. These are called ANHYDRIDES.

Oxides of nonmetallic elements like carbon, sulfur, and nitrogen produce an acid in water.

CO2 (oxide) + H2O ( H2CO3

Oxides of metallic elements like sodium, magnesium, and calcium usually form bases.

CaO (oxide) + H2O ( Ca+2 + 2OH-

PRACTICE:

1. NO2 + H2O (

2. SO3 + H2O (

3. MgO + H2O (

NAMING ACIDS

Binary Acids

Binary Acids contain only two elements. They have the prefix hydro-and the suffix -ic . Find the stem of the element that combines with hydrogen.

HCl Hydro- chlor -ic acid Exception: HCN Hydrocyanic acid

HBr Hydro- brom -ic acid

Ternary or Oxy Acids

These contain hydrogen, oxygen, and one other element. The stem is determined by what element is combined with oxygen and hydrogen in the acid. Suffix is determined by the number of oxygen atoms.

Generally, the ion with a suffic of –ate become acids with the suffix, -ic.

H2SO4 Sulfuric acid from sulfate

HNO3 Nitric acid from nitrate

H3PO4 Phosphoric acid from phosphate

HClO3 Chloric acid from chlorate

Generally, the ion with a suffix of –ite become acids with the suffix, -ous.

H2SO3 Sulfurous acid from sulfite

HNO2 Nitrous acid from nitrite

HClO2 Chlorous acid from chlorite

Carboxylic Acids

These are organic acids. The one we will study is acetic acid, HC2H3O2.

NAMING BASES

To name a base, you just follow the rules learned for naming any ionic compound.

VIDEODISC: Acid Strength vs. Concentration

1. What is the difference between the concentration of an acid and its strength?

2. 0.1M HCl has pH of 1.0; 0.10M H2SO4 has a pH of 1.0 and 0.1M acetic acid has a pH of 3.0. Which would you say is the weakest?

STRONG VS WEAK, CONCENTRATED VS DILUTE

These terms are not interchangeable. Strong vs. weak deal with the amount of ionization. Concentrated vs. dilute deal with molarity or some other form of concentration.

STRENGTHS OF ACIDS

Which is stronger, HNO3 or H3PO4?

Can you rank the strength of HF, H3BO3, and H2SO3 based on the formula alone?

DEMO: HCl and Acetic acid

1. What happens when the molarity of HCl and HC2H3O2 goes down?

2. Why does HCl have a brighter glow than HC2H3O2?

Ions can carry electricity through a solution and HCl forms more ions than acetic acid does. A strong acid readily transfers H+ ions to water to form H3O+ ions. Strong acids react completely to form ions and are therefore, strong electrolytes. These acids produce the maximum amount of ions.

HCl (aq) + H2O(aq) ( H3O+(aq) + Cl-(aq)

strong acid - one way arrow

A weak acid transfers only some of its H+ ions to water. Weak acids, because so many fewer ions are formed, are weak electrolytes. The ions only partially ionize.

One way to signify a strong acid from a weak acid is by their equations:

HC2H3O2(aq) + H2O(aq) ↔ H3O+(aq) + C2H3O2-(aq)

weak acid - reversible arrow

Strong Acids Weak Acids

HCl H2SO4 HC2H3O2 HClO

HBr HClO4 HCN HCO3-

HI HNO3 HF HNO2

The reason behind this has to do with acid strength and the Brønsted-Lowery Model.

HX + H2O ( H3O+ + X-

acid base CA CB

Water is a stronger base (in the forward reaction) than is the conjugate base, X-. The ionization equilibrium lies almost completely to the right because water has a much greater attraction for H+ than does the base X-. It is basically a battle of the bases.

The Brønsted-Lowery Model helps explain acid strength, but the model does not provide a quantitative way to express the strength of an acid or to compare the strengths of various acids. The equilibrium constant expression provides the quantitative measure of acid strength.

STRENGTHS OF BASES

The substances that have a strong affinity for H+ are called strong bases. Strong bases are also strong electrolytes because they completely ionize in water.

Weak bases only partially ionize and are considered weak electrolytes.

Strong Bases: Weak Bases:

CaO NH3

NaOH H2NNH2

KOH CO3-2

Ca(OH) 2 PO4-3

Some metallic hydroxides, like Ca(OH)2 and Mg(OH)2 have low solubility and are poor sources of OH-. But they are considered strong bases because what does dissolve ionizes completely.

IONIZATION CONSTANT

For any given reaction of an acid and water, an equilibrium expression can be written:

HCN (aq) + H2O (l) ↔ CN- (aq) + H3O+ (aq) Keq = [CN-][H3O+]

[HCN][H2O]

Water is at a constant concentration. So it moves to the left side. The left side, Keq[H2O], now becomes Ka, the ionization constant.

Keq[H2O] = [CN-][H3O+] Ka = [CN-][H3O+]

[HCN] [HCN]

Ka is the value for the ionization of an acid. Ka can tell you if reactants or products are favored at equilibrium. See page 647 for some weak acid values.

If Ka > 1, the acid is strong. (more product)

If Ka < 1, the acid is weak. (more reactant)

Weak bases can be written the same way, except the ionization constant is Kb. See page 649.

PRACTICE:

1. Write the acid ionization equation and the ionization constant expression for HCOOH.

2. Write the base ionization equation and the ionization constant expression for NH3.

pH Scale

The ionization of water into H3O + ions is very slight. But knowledge of this helped scientists build an acidity scale.

The pH scale is the measure of hydronium ion concentration, and thus, acidity.

The concentration of H3O+ is expressed in powers of 10, from 1014 to 100. For convenience, scientists use:

pH = -log[H3O+] We use the negative log so pH is always positive.

Examples:

[H3O+] = 1 x 10-3 [H3O+] = 6.3 x 10-8

pH = -log[1 x 10-3] pH = -log[6.3 x 10-8]

pH = 3.0 pH = 7.20

The smaller the pH, the more acidic the solution.

A solution with a [H3O+] of 10-1M - a strongly acidic solution - has a pH of 1.

A solution with a [H3O+] of 10-7M - a neutral solution - has a pH of 7.

A solution with a [H3O+] of 10-13M - a strongly basic solution - has a pH of 13.

Remember that the pH scale is logarithmic. Each one-unit change on the scale represents a 10-fold change in the concentration of H3O+ ions.

As the hydronium ion concentration increases, does pH increase or decrease?

Significant Figures

For any logarithm, the number of digits after the decimal point should equal the number of significant digits in the original number. That means that the total number of significant digits in the concentration should equal the number of digits past the decimal point in the pH number.

PRACTICE:

1. [H3O+] = 7.3 x 10-5 M 3. [H3O+] = 5.0 x 10-2 M

What is the pH? What is the pH?

2. [H3O+] = 6.23 x 10-4 M 4. [H3O+] = ? M

What is the pH? pH = 3.000

Going Backwards:

If given pH, you can calculate [H3O+].

Example: Find the concentration of a solution with a pH of 9.0

pH = -log [H3O+] antilog (-9.0) = [H3O+] 1 x 10-9 = [H3O+]

9.0 = -log[H3O+] -9.0 = log[H3O+]

PRACTICE

1. The pH is 3.57. 2. The pH is 12.32.

What is the [H3O+]? What is the [H3O+]?

Other ways to calculate pH:

If given [OH-], you can calculate pH. [OH-] [H3O+] = 1 x 10-14

Example: Find the pH of a solution with an [OH-] of 1.00 x 10-9 M.

[OH-][H3O+] = 1 x 10-14 [H3O+] = 1.00 x 10-5

pH = -log[1.00 x 10-5]

[H3O+] = 1 x 10-14 = 1 x 10-14 pH = 5.000

[OH-] [1.00 x 10-9]

1. [OH-] = 3.5 x 10-5 M. 2. [OH-] = 9.65 x 10-11 M

Find the pH. Find the pH

Going Backwards again:

pOH is a scale to measure alkalinity. pH + pOH = 14

pOH = -log[OH-]

Example: Find the pOH of an [OH-] of 1 x 10-9

pOH = -log[OH-]

pOH = -log[1 x 10-9]

pOH = 9.0

PRACTICE:

1. The pOH is 3.57. **2. The pOH is 12.32.

What is the [OH-]? What is the [H3O+]?

]]]

MEASURING pH]]]

Three methods:

1. Litmus paper, indicators - usually measure a range and are not as accurate.

2. pH paper – measure within a number or two and more accurate than litmus.

3. pH meter – electronically measures pH

NEUTRALIZATION REACTIONS

The reaction of an acid and a base to form a salt and water.

HClO3(aq) + NaOH(aq) ( H2O(l) + NaClO3(aq)

Acid base water salt

Not all neutralization reactions form water. If the base is a hydroxide, it will. A reaction between a basic anhydride and an acid anhydride just forms a salt.

Not all neutralization reactions form a neutral solution. It is possible to get acidic or basic salts, if the salt still has an ionizable hydrogen atom.

TITRATION

To measure the concentration of an acid or base in solution, a titration can be performed.

A TITRATION is a carefully controlled neutralization reaction. A standard solution is used to determine the concentration of another solution. The concentration of the standard solution is known.

By adding an indicator, the point at which complete neutralization occurs can be determined.

To perform a titration, follow these steps:

1. Measure a standard volume of the acid or base with the known concentration.

2. Fill the buret with the solution of unknown concentration.

3. Slowly add the unknown and mix until the equivalence point is met.

The equivalence point is where the moles of H+ equal the moles of OH- in both solutions.

NaOH + HCl ( NaCl + H2O

Na+ + OH- + H+ + Cl- ( Na++ Cl- + H2O

1 mole H+ and 1 mole OH-

A chemical dye called an indicator is used instead of a pH meter. The equivalence point, or end point, is where the indicator changes color.

CALCULATING MOLARITY

1. Write the balanced chemical equation

2. Find the moles of the standard solution.

*Remembering the molarity = mol solute divided by liters of solution, solve for moles of solute.

3. Use the mole ratio to determine the number of moles of the unknown.

4. Calculate the molarity of the unknown using the moles from #3 and the volume titrated.

VIDEODISC: Rainbow Indicators

1. What is a universal indicator?

2. So where did the stain go? (Squirt bottle contains dilute ammonia.)

3. What indicator was used to treat his shirt?

INDICATORS

An indicator is a weak acid or base that undergoes a color change when it gains or loses an H+ ion.

How They Work:

HIn + H2O ↔ H3O+ + In-

When you add an acid, you add H3O+ ions. This shifts the equilibrium to the left, increasing [HIn]. In litmus, HIn is colored red. Adding a base, which removes H3O+ ions, shifts the equilibrium to the right, increasing [In-]. For litmus, In- is colored blue.

Each indicator has its own equilibrium constant and thus a different color range for different pH. See table, page 662.

Because they are viewed by sight and thus subjective, they cannot be used to measure precise pH changes.

Examples:

litmus phenolphthalein alizarin yellow

methyl orange thymol blue alizarin

methyl red bromothymol blue universal

methyl violet Bromophenol blue

BUFFERS

pH must, at times, be controlled within very narrow limits. One way to do this is with a buffer. A BUFFER is a solution that is able to release or absorb H+ ions, keeping a solution’s pH constant. It resists a change in pH. The mixture of ions and molecules in a buffer solution resists change in pH by reacting with any H+ or OH- added to the solution. The most common buffers are mixtures of weak acids and their conjugate bases or vice versa.

How They Work:

Consider a buffer of acetic acid and the acetate ion:

H3O+ + C2H3O2- ↔ H2O + HC2H3O2

Adding an acid [H3O+] reacts with the acetate ion to form acetic acid. Almost all react and it forms mostly product. Because very little are left in solution, pH changes only slightly. The same is true if a base is added:

OH- + HC2H3O2 ↔ H2O + C2H3O2-

All buffers have a limited capacity to neutralize added H3O+ ions or OH- ions. The amount of acid or base that a buffer can neutralize is called the buffer capacity. The greater the concentration of buffering molecules, the greater the buffering capacity.

Buffers in Blood

A pH of 7.4 must be maintained in the bloodstream. It can deviate only within narrow limits – 0.3 pH units up or down. Acidosis occurs if the blood pH falls below 7.1. Alkalosis occurs when the blood pH rises above 7.7 (Some experience cramping due to lactic acid build-up in muscles.) The body maintains pH in a variety of ways:

1. Excess acids and bases are excreted in the urine.

2. Carbon dioxide is eliminated by breathing.

3. The buffers H2CO3 and HCO3- are used. The reaction is as follows: CO2 + H2O ↔H2CO3 ↔ H+ + HCO3-

a. If the acid level rises, equilibrium shifts left to consume H+ and we breathe more to get rid of the excess carbon dioxide.

b. If the base level rises, equilibrium shifts right and breathing slows down to slow down the removal of carbon dioxide. (This is also what we are doing when we ask hyperventilators to breathe in a bag!)

VIDEODISC: Does Aspirin Buffer?

1. After the buffered aspirin dissolves in beakers three and four, the color changes from clear to pink. What do you conclude from this?

2. The color in beaker three did not change, but the color in beaker four did. What is going on?

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