ACIDS and BASES

Ionic solids which ______________________________ are not the only substances which dissociate into ions.

Such solids or "salts" may be considered to be the ______________________________ between two other classes of compounds which also exhibit dissociation behavior in water.

These compounds are known as acids and bases.

Like the solids which they form when they react, some dissociate 100% and some, like the insoluble salts, (which dissociate to a very small extent - CaCO₃). Therefore, some solids may be treated simply (the STRONG acids and bases) - and others require treatment as equillllibrium systems (the WEAK acids and bases).

Acids and bases are common substances which we encounter in everyday life. They have been known for many years and the original descriptive definitions which have come down to us are still useful and interesting to look at:

ACIDS

conduct electricity
taste sour
react with "active" metals to produce H_{2 (g)} [Refer to the Activity Series of Metals in the BALANCING EQUATIONS & CHEMICAL REACTION TYPES section]
neutralize bases
change litmus to ______________________________

* Indicators: An indicator is a substance that is one color in an acidic solution and another color in a basic solution.

common examples:
citric acid
acetic acid
hydrochloric acid
sulfuric acid

BASES

conduct electricity
______________________________
feel slippery
neutralize acids
change litmus to ______________________________

common examples:
sodium hydroxide
potassium hydroxide
ammonia

ACID / BASE THEORIES:

In addition to the operational definitions for acids and bases, a number of theories have been developed to explain the behavior of these substances.

One of earliest and simplest is the Arrhenius Theory. In it, acids and bases are defined in terms of their behavior ______________________________.

Arrhenius Acid: dissolves in water to form H⁺¹ ions (commonly in the form of hydronium ion, H₃O⁺¹).
Arrhenius Base: dissolves in water to form OH^-1 ions.

For example:

HCl_(l) ® H⁺¹_(aq) + Cl^-1_(aq)

HCl is an acid

NaOH_(s) ® Na⁺¹_(aq) + OH^-1_(aq)

NaOH is a base

The careful construction of the definitions will be appreciated when we consider the base ammonia, NH₃:

NH_3(g) + H₂O_(l) ® NH₄⁺¹_(aq) + OH^-1_(aq)

The water is required here to make sense out of the basic (or alkaline) behavior of ammonia in water, even though it does not contain hydroxide ions in its structure.

Because it is so simple to use, we will frequently fall back on the Arrhenius theory.

In the 1920's the Arrhenius Theory was supplanted by a theory developed simultaneously by Brønsted and Lowry. This model describes the reactions between acids and bases in terms of a competition for protons.

Brønsted-Lowry Acid: a PROTON ______________________________ (any compound that donates a H⁺¹)
Brønsted-Lowry Base: a PROTON ______________________________ (any compound that accepts a H⁺¹)

In practice, the results are the same as in the Arrhenius system, although a few more substances may be classified as acids or bases, and ammonia, NH₃, does not require any special treatment.

In terms of ordinary use, the Brønsted-Lowry Theory contributes 2 important concepts to our understanding of acids and bases:

1) H⁺¹ is unlikely to exist in water because of its high charge density; therefore H₃O⁺¹ is preferred

2) all acid/base behavior may be considered in terms of reactions between relative acids and bases to produce weaker relative acids and bases (conjugate acid/base pairs).

* * By comparison, a weak acid has a strong conjugate base, and a strong acid has a weak conjugate base.

The second concept could use a few examples for explanation:

· The pairs shown are called conjugate pairs (the conjugate acids and bases are the ______________________________ ) and differ only by one proton (H⁺¹).

· Substances such as water, have the ability to act as either a base or an acid. [H₃O⁺¹ or OH^-1 ]

In the pair, the one with the extra proton is the acid.

The relative nature of acids and bases can be shown by selecting a reaction in which water behaves as a base:

Practice Problems:

Determine the Brønsted-Lowry Acid, Brønsted-Lowry Base, conjugate acid and conjugate base for the following reactions:

1. H₃PO₄ + HOH ® H₂PO₄^-1 + ₃O⁺¹

2. NH₃ + HOH ® NH₄⁺¹ + OH^-1

ANSWERS:

HOMEWORK PROBLEMS:

1a. Get worksheet from Mr. Craig - "Brønsted-Lowry Acids and Bases"

ANSWERS:

Practice Problems:

The following substance acts as a Brønsted-Lowry Acid in water. Write the chemical equation that illustrates its reaction with water.

3. HCO₃^-1

The following substance acts as a Brønsted-Lowry Base in water. Write the chemical equation that illustrates its reaction with water.

4. HCO₃^-1

ANSWERS:

HOMEWORK QUESTIONS:

2a. The following substances act as Brønsted-Lowry acids in water. Write the chemical equation for each that illustrates its reaction with water.

i. hydroiodic acid, HI

ii. ammoium ion, NH₄⁺¹

iii. H₂CO₃

iv. HNO₃

2b. The following substances act as Brønsted-Lowry bases in water. Write the chemical equation for each that illustrates its reaction with water.

i. cynide ion, CN^-1

ii. SO₄^-2

iii. C₂H₃O₂^-1

iv. NH₃

ANSWERS:

STRONG ACIDS and BASES:

The ______________________________ of an acid or base (as opposed to its concentration) is a measure of how much it dissociates in water. The so-called ______________________________ acids and bases dissociate essentially 100% and thus need not be treated as equilibrium systems in themselves.

However, since dissociation takes place in water, an understanding of the equilibrium that exists in pure water and all water solutions is helpful.

Liquid water dissociates to a very small extent at room temperature:

H₂O_(l) ® H⁺¹_(aq) + OH^-1_(aq)

This equilibrium lies far to the left, and in fact the value for K_eq at 25⁰C is 1.00 x 10^-14 . Since this is such an important constant, it is given its own symbol: K_w. It has a denominator of 1 since the reactant is a pure liquid. Thus K_w is:

[H⁺¹][OH^-1] = 1.00 x 10^-14

Obviously in "neutral" or pure water the concentrations of hydrogen and hydroxide ions are equal at 1.00 x 10^-7 M.

In an acid solution [H⁺¹] > [OH^-1]

In a base solution [H⁺¹] < [OH^-1]

O.K., so what does this have to do with strong acids and bases?

Recall the definitions (Arrhenius will do). Strong acids will ______________________________ solvent with H⁺¹ ions.

This forces the water equilibrium to the left, reducing the number of hydroxide ions. The converse can be said for the addition of a strong base to water.

But whatever the case, the product of hydrogen and hydroxide ion concentrations must equal 1.00 x 10^-14 (at 25⁰C, of course!). Thus we have a way to calculate the concentration of both ions in an solution of a strong acid or base.

examples:

What are both H⁺¹ and OH^-1 ion concentrations in a 0.5 M solution of nitric acid?

[H⁺¹] = 0.5 M since nitric acid is STRONG and dissociates 100% (hydrogen ion from water is negligible)

[OH^-1] = (1.00 x 10^-l4) / 0.5 = 2 x 10^-14 M

What is [OH^-1] in a 0.1 M HC1 solution?

since [H⁺¹] = 0.1 M (HCl is STRONG), then:

[OH^-1] = (1.00 x 10^-14) / 0.1 = 1 x 10^-13 M

What is [H⁺¹] in a 0.01 M KOH solution?

since [OH^-1] = 0.01 M (KOH is STRONG), then:

[H⁺¹] = (1.00 x 10^-14) / 0.01 = 1 x 10^-12 M

A good question comes to mind: how can you tell if an acid or base is strong? For bases the answer is easy: strong bases contain metal ions from Groups I or II.

WEAK ACIDS and BASES

The WEAK acids and bases are those which are not ______________________________, i.e., those which ______________________________ dissociate to any great extent in water solutions.

As such, they must be considered as equilibrium systems and their quantitative treatment is more complex than that of their strong counterparts.

A typical example of a weak acid is acetic acid:

HC₂H₃O_2(aq) ® H⁺¹_(aq) + C₂H₃O₂^-1_(aq)

Therefore, at equilibrium a considerable amount of HC₂H₃O₂ molecules remain.

Practice Problems:

5. Calculate the [OH^-1] if the [H⁺¹] = 7.45 x 10^-6 M. Then indicate if the solution is acidic or basic.

6. Calculate the [H⁺¹] if the [OH^-1] = 1.87 x 10^-3 M. Then indicate if the solution is acidic or basic.

ANSWERS:

HOMEWORK PROBLEMS: Note: [H⁺¹][OH^-1] = 1.00 x 10^-14 M

3a. Calculate the [OH^-1] if the [H⁺¹] = 1.00 x 10^-4 M. Then indicate if the solution is acidic or basic.

3b. Calculate the [H⁺¹] if the [OH^-1] = 4.50 x 10^-9 M. Then indicate if the solution is acidic or basic.

3c. Calculate the [H⁺¹] if the [OH^-1] = 8.30 x 10^-2 M. Then indicate if the solution is acidic or basic.

ANSWERS:

What you have learned about acids and bases so far has provided you with the essential tools to begin investigating how acids and bases react with each other.

You learned the Brønsted-Lowry definitions of acids and bases.

Exactly what happens when an acid and a base are mixed together? How could you tell that a chemical reaction has taken place? To answer these questions we will discuss the reaction between two common substances found in the laboratory, HCl (hydrochloric acid) and NaOH (sodium hydroxide).

NEUTRALIZATION - REACTIONS OF ACIDS and BASES:

The reaction of an acid with a base---both of which dissociate in water---produces another substance which dissociates in water, and WATER itself!

These products of the neutralization of acids with bases are known collectively as ______________________________.

A salt is simply an ionic compound which dissociates in water to some extent. Generally, salts are composed of metal cations (the cation comes from the base) and non-metal anions (the anion comes from the acid).

There are, however, a few exceptions, the most important of which for us is "ammonium, NH₄⁺¹."

So the ______________________________ might be expressed:

This is a double displacement reaction in which the first part of the acid and the last part of the base combine to make HOH or water, and the other parts form some salt.

Of course, if the salt is soluble, it will not form a solid unless the water is removed (evaporated). Therefore, many acid/base reactions simply result in the formation of water (the net-ionic equations would all be the same). Only when the salt is ______________________________ will it show up in the reaction.

The one notable exception to this general scheme is reactions with the base ammonia, NH₃. Since ammonia contains no hydroxide ion, reactions with acids produce no water and do not neatly fit the four categories of reactions we have studied.

EVIDENCE OF AN EXOTHERMIC REACTION:

In this case, one indication of a reaction can be found by feeling the flasks before and after the reaction. If ______________________________, the solution is distinctly warmer after mixing. However, this change in ______________________________ does not mean that a chemical change took place - a physical change could have occurred.

DEFINING NEUTRALIZATION:

Acids and Bases can be considered opposites because their behaviors can be described as either proton donors or proton acceptors. Chemists say that an acid ______________________________ a base when the two react completely, leaving no excess acid or base.

Similarly, when sodium hydroxide neutralizes hydrochloric acid, the resulting solution has neither acidic nor basic properties. Reactions between an acid and base may be classified as a ______________________________.

Neutralization may occur between acids and bases other than HCl and NaOH.

COMPLETE NEUTRALIZATION:

The salt produced depends on whether the acid reacts with the base in a one-to-one ratio or a one-to-two ratio. Acids can donate more than one proton are called ______________________________ acids.

example, phosphoric acid, H₃PO₄:

For a reaction in which all of the acidic protons (hydrogen ions, H⁺¹) leave the acid, the term complete neutralization is used.

TITRATION:

A frequently used technique for comparing the concentrations of two solutions is known as titration.

A measured sample of ______________________________ concentration is "titrated" with an acid or base of known concentration-- generally added from a buret--until the added H⁺¹ or OH^-1 has been exactly neutralized.

At that point, the pH of the solution may be 7, or less, or more, depending on the salt formed during neutralization. In any case, an indicator may be used to show the "endpoint" of the titration, or instruments may detect the sudden change in pH as the equivalence point is passed.

It is important to realize that at the equivalence point, the moles of hydrogen ion added equals the moles of hydroxide ion neutralized, or visa versa. It is not always true, however, that the concentrations of these two ions are equal. If that were true every equivalence point would be at pH 7.

The use of an ______________________________ helps to determine when a certain pH has been achieved (which will be demonstrated in class). Typically, an indicator is placed in an acidic solution and a base is added to the acid (during a titration). When the desired pH has been reached the indicator will change the clear solution (acid mixed with base) into a colored (not cloudy) solution.

Since the moles of hydrogen and hydroxide ions are equal, the calculation of either can be achieved with a formula which equates moles of two mixed solutions:

As before, in the dilution formula, the volume units do not have to be in liters, but they must match. Here, however, the second volume is NOT the total, but the volume of the base.

example

If it takes 50 mL of 0.30M HCl to titrate 100 mL of a NaOH solution, what is M_b?

(50 mL)(0.30M) = (100 mL)M_b
M_b = 0.15M

Titration becomes a little confusing with polyprotic acids and their base counterparts, since one molecule may bring 2 or more hydrogen or hydroxide ions with it. An adjustment must be made in the molarity values in these cases.

example

What is the M of an Al(OH)₃ solution if 10.5 mL is neutralized by 15.0 mL of 0.20M H₂SO₄?

the acid molarity must include a factor of 2 for the two hydrogens in the acid; similarly, the base molarity must include a factor of 3:

(15.0 mL)(2 x 0.20M) = (10.5 mL)(3 x M_b)

thus, M_b = 0.19M

Also, and you will learn this in AP Chemistry, stoichiometry MAY BE used (for ALL titration problems) and never have to worry about a "multiplication factor"...

Practice Problem - Let's do these BOTH ways - using M_aVa and using stoichiometry...

7. When 42.5 mL of 1.03 M NaOH is added to 50.0 mL of vinegar (a solution of acetic acid, HC₂H₃O₂), the phenolphthalein in the solution just turns pink. Calculate the concentration of acetic acid in vinegar.

8. Calculate the volume of a 2.55 M HNO₃, an acid, needed to neutralize 67 mL of barium hydroxide, a base.

ANSWERS:

HOMEWORK QUESTIONS:

4a. In a titration of another sample of vinegar, you find that it requires 11.10 mL of 0.748 M NaOH to neutralize a 10.0 mL sample of vinegar. What is the concentration of acetic acid in the sample of vinegar?

4b. What is the concentration of acid in rainwater when 100.0 mL is titrated with 25.12 mL of 0.00105 M NaOH? Since acid rain contains several acids, use the symbol HA to represent the acids that are present.

4c. What does an indicator "indicate"?

ANSWERS:

pH SCALE:

The pH scale is an easy way to classify how strong an acid or a base is in a solution.

The pH scale is a scale from 0 to 14. On the scale, 0 is the strongest acid, 14 is the strongest base, and a neutral solution is located at 7.

So if you know that substance had a pH of 6, you would know that it was a weak acid. If you were given a substance with a pH of 13, you would know that it was a strong base. Also, if you were given the H⁺¹ concentration of an acid or base you can use the below formula to find the pH for that substance.

If you were given the OH^-1 concentration of an acid or base, you would have to use the below equation.

The extent to which an acid or base dissociates in water will determine the relative acidity or alkalinity of the solution. Chemists have found it convenient to express this property as a power of ten. For example, in a neutral solution, [H⁺¹] = 1 x 10^-7M. Thus the "pH" of that solution is 7.

A "p" function in chemistry is a negative log (base 10) function. And so pH is defined as -log [H⁺¹] (or, alternatively, log 1 / [H⁺¹]).

A companion function known as pOH similarly expresses the hydroxide concentration. Although pH is more common, both are used. They are related simply:

pH + pOH = 14.

To summarize:

examples

a. Find pH if [H⁺¹] = 3.1 x 10^-3M

b. What is the pOH?

Since pH + pOH = 14, then:

c. If pH = 5.7, what is [H⁺¹]?

[emphasize the order of things here--the sign must be changed first and then the anti-log taken]

Practice Problems:

9. Calculate the pH of a 0.01 M nitric acid solution, HNO₃? Nitric acid is a strong acid.

10. Calculate the pOH of 2.3 x 10^-4 M sodium hydroxide solution.

11. Calculate the pH of a 0.0045 M lithium hydroxide solution.

ANSWERS:

HOMEWORK QUESTIONS: Recall: [ H ⁺¹ ] = [ H₃O ⁺¹ ]

5a. What is the pH of a 0.50 M solution of HCl? HCl is a strong acid.

5b. Find the pH of a solution whose H₃O⁺¹ concentration is
i. 0.1 M

ii. 1.0 x 10^-7 M

iii. 2.55 x 10^-3 M

ANSWERS:

12. What is the H₃O⁺¹ concentration of a solution that has a pH of 3.00 ?

13. Calculate the [OH^-1] for a solution that has a pOH of 7.55.

14. Calculate the [H⁺¹] for a solution that has a pOH of 2.45.

ANSWERS:

6a. Find the value of [H₃O⁺¹] in a solution that has a pH of 2.

6b. Find the value of [H₃O⁺¹] in a solution that has a pH of 13.

6c. Find the value of the [OH^-1] in a solution with a pH of 3.75.

ANSWERS:

COMMON SOLUTIONS:

Some examples of common solutions and their pH are given in the table below: