11. introduction to acids, bases, ph, and · pdf fileintroduction to acids, bases, ph, ... the...

C. Graham Brittain of 14 11/23/2010

11. Introduction to Acids, Bases, pH, and Buffers

What you will accomplish in this experiment

You’ll use an acid-base indicating paper to:

• Determine the acidity or basicity of some common household substances

• Compare the pH of two different concentrations of strong acid and strong base solutions

• Compare the pH of strong and weak acid solutions that have equal concentrations

• Compare the pH of strong and weak base solutions that have equal concentrations

• Investigate the impact of adding a strong acid or strong base to a buffer system.

Concepts you need to know to be prepared

Bronsted-Lowry Acids and Bases

In 1923, Johannes Bronsted of Denmark and Thomas Lowry of England independently proposed the same theory

about acid and base compounds and the chemical reaction that occurs between them. The Bronsted-Lowry Theory

focuses on the movement of the hydrogen ion, H+ (which is usually referred to as just a “proton”).

Bronsted and Lowry define an acid as a compound which can donate a proton. They define a base as a compound

that can accept a proton. An acid-base reaction is then simply the transfer of a proton from the acid to the base.

Even the simple dissolution of an acid in water results in a Bronsted-Lowry acid-base reaction, because the acid

molecule donates a proton to the water molecule.

For example, when the molecular compound hydrogen chloride dissolves in water, it “dissociates” into ions,

producing a proton, H+, and a chloride ion, Cl

-. The dissociated proton is donated to a water molecule to produce

the hydronium ion, H3O+.

HCl (aq) + H2O (l) � H3O+

(aq) + Cl−

(aq)

As the HCl molecule is donating the proton, it is acting as a Bronsted-Lowry acid. For this reason, we refer to

solutions of HCl in water as “hydrochloric acid.”

As the H2O molecule is accepting the donated proton, it is acting as a Bronsted-Lowry base.


According to Bronsted-Lowry theory, every acid-base reaction creates a “conjugate acid-base pair.”

In the acid-base reaction shown below, nitric acid (HNO3) donates a proton (H+) to a water molecule. The

polyatomic nitrate ion (NO3-) is said to be the “conjugate base” of nitric acid.

Similarly, the hydronium ion, (H3O+) is called the “conjugate acid” of the H2O molecule. Thus HNO3 and NO3

-

are a “conjugate acid-base pair;” H2O and H3O+ are also a “conjugate acid-base pair.”

“Strong” and “Weak” Acids

This idea of conjugate acid-base pairs raises an interesting question:

If the hydronium ion (H3O+) is truly an acid, and can donate a proton to the nitrate ion base (NO3

-), why doesn’t the

above reaction reverse itself? Why don’t hydronium ions (H3O+) donate their protons (H

+) back to the nitrate ions

(NO3-) to re-form the HNO3 and H2O molecules?

The answer is that not all acids have the same proton-donating capabilities, and not all bases are equal in their

proton-accepting capabilities.

Acids that are particularly good proton donors (dissociating nearly 100% to produce large amounts of hydronium

ion) are said to be “strong acids.” Conversely, good proton acceptors (generating large amounts of hydroxide

ion) are called “strong bases.”

In the above reaction, HNO3 is such an exceptionally strong acid that the dissociation (ionization) reaction of

hydrogen chloride with water goes essentially to completion.

And as the hydronium ion, H3O+, is a comparatively weak acid (and the NO3

-, is an extremely weak base: a

miserably poor proton acceptor), the above reaction has no tendency at all to reverse itself.

The number of acids that are considered to be “strong” is actually quite small; your lecture textbook lists just a

few “strong” acids:

• Hydrochloric acid (HCl)

• Hydrobromic acid (HBr)

• Nitric Acid (HNO3)

• Sulfuric Acid (H2SO4)

All other acids are considered to be “weak,” in that they vary considerably in the extent to which they “dissociate”

in water (donate their protons to water molecules) to produce the hydronium ion, H3O+. Because the dissociation

of weak acids is an equilibrium process, we can use equilibrium constants to tell us quantitatively just how

“strong” any weak acid is; that is, how readily the weak acid donates its proton to water.

An equilibrium constant describes the “position” of the equilibrium by expressing the concentrations of the

products and reactants that are present in solution as a ratio:

� = ��

��

The brackets [ ] mean “concentration of,” and the unit used to express the concentration is always molarity (moles

of solute per liter of solution, or moles/L, M).


The dissociation of ANY weak acid (HA) in water can be represented as:

HA (aq) + H2O (l) � H3O+

(aq) + A -

(aq)

And the equilibrium constant expression (K) for the dissociation of ANY weak acid can be written as:

�� = ��

��

As this equilibrium constant describes the extent of dissociation of a weak acid in water, it’s known as the Acid

Dissociation Constant, Ka. The Ka values of several weak acids are shown below.

The Amphiprotic Water Molecule: Self-Ionization of Water

Water molecules are remarkable in that they can undergo “self-ionization:”

• One water molecule acts as an acid and donates a proton;

• A second water molecule behaves as a base and accepts the proton.

The products of this Bronsted-Lowry acid-base reaction are the hydroxide ion, OH-, and the hydronium ion, H3O

+.

A compound that can act as either an acid or a base is said to be “amphiprotic.”

H2O (l) + H2O (l) � H3O +

(aq) + OH -

(aq)

The above reaction occurs in pure water to only a very small equilibrium extent, that is, very few hydroxide and

hydronium ions are formed.

When the self-ionization reaction reaches equilibrium (and the rates of the forward and reverse reaction are equal),

only 1.0 x 10-7

moles of hydroxide ion and 1.0 x 10-7

moles of hydronium ion can be found in one liter of water.


These concentrations can be expressed mathematically as follows:

[H3O+] = [OH‾] = 1.0 x 10

−7 M (moles/liter)

And the extent of the equilibrium can be described by writing an equilibrium constant expression called the Ion

Product of Water, Kw:

�� = �� = �1.0 × 10��1.0 × 10�� = 1.0 × 10� !

The value of the ion product of water applies not only to pure water, but also to any aqueous solution. This is very

convenient, because if we know the concentration of ONE of these two ions, we can calculate the concentration of

the OTHER.

For example, if a solution of hydrochloric acid has a concentration of 0.010 M (0.010 moles of HCl per liter), the

hydronium ion concentration [H3O+] = 0.010 M = 1.0 x 10

-2 M (because the strong acid HCl dissociates

completely to produce the hydronium and chloride ions).

From the ion product of water, the hydroxide ion concentration [OH‾] in this HCl solution can be calculated as:

�� = ��

�� =

1.0 × 10� !

1.0 × 10�"= 1.0 × 10� " #

As you can see from this example, the HCl solution has a higher concentration of hydronium ions than does pure

water (1.0 x 10-2

M in HCl, only 1.0 x 10−7

M in water), and a lower concentration of hydroxide ions than does

pure water (1.0 x 10-12

M in HCl, but 1.0 x 10−7

M in water).

The concentrations of hydronium and hydroxide ions in any aqueous solution are always related such that the

product of their concentrations is equal to the value 1.0 x 10−14

(Kw, the Ion Product of Water).

The pH Scale

Soren P. L. Sorenson introduced the pH scale in the early 1900s as a simplified way of expressing the

concentration of hydronium ions in an aqueous acid or base solution.

The pH of a solution is defined as the negative logarithm of the hydronium ion concentration:

�� = −%�&��

You should recall that in pure water, the hydronium ion concentration is 1.0 x 10−7

M. The logarithm of this

concentration is -7 (the value of the exponent), thus the pH of pure water is equal to 7 (the negative of the

exponent’s value).

Because water is amphiprotic (can act as either an acid or a base), and the concentrations of the hydronium and

hydroxide ions are equal in pure water, pH 7 is defined as a “neutral” solution.

The 0.010 M of the hydrochloric acid solution has a hydronium ion concentration equal to the molarity of the

strong acid: [H3O+] = 0.010 M = 1.0 x 10

-2 M (because the strong acid dissociates 100%). The logarithm of this

value is -2, so the pH of this HCl solution is equal to 2.

On the pH scale, acidic solutions are those that have a pH less than 7. This is because acids donate protons to

water, causing the concentration of H3O+ to become greater than that in pure water (greater than 1.0 x 10

−7 M, so a

less negative exponent), and thus the pH of the solution decreases from the neutral value of 7 to a smaller number.

Basic solutions are those that have a pH greater than 7. This is because bases either release hydroxide ions (OH-)

as they dissolve in water, or bases accept protons from water molecules to create hydroxide ions.

In either case, because the concentrations of the hydroxide and hydronium ions are always related by the ion

product of water (Kw = [H3O+] [OH‾] = 1.0 x 10

-14), as the concentration of the OH

- ion increases (becomes greater

than that in pure water), the concentration of H3O+ ion decreases (becomes less than 1.0 x 10

−7 M, a more negative

exponent). Thus the pH of the basic solution increases from the neutral value of 7 to a larger number.


“Strong” and “Weak” Bases

Two examples of bases are presented below, and like acids, they are commonly classified as “strong” or “weak,”

depending on their ability to generate hydroxide ions.

“Strong” bases are typically ionic compounds composed of a hydroxide ion, OH-, and a Group I Alkali Metal ion

(such as sodium or potassium). These salts are extremely water soluble, thus they dissociate completely to produce

a solution of hydroxide ions.

NaOH (s) →OH2

Na+

(aq) + OH-

(aq)

“Weak” bases can include less water-soluble hydroxide salts, such as Ca(OH)2, or molecular compounds such as

NH3.

The molecular compounds act as classic Bronsted-Lowry bases in that they accept a proton from a water molecule

to yield aqueous hydroxide ions. This reaction is reversible, and usually occurs to only a small equilibrium extent.

NH3 (aq) + H2O (l) � NH4+

(aq) + OH-

(aq)

The dissociation of ANY weak base (B) in water can be represented as:

B (aq) + H2O (l) � OH-

(aq) + HB +

(aq)

And the equilibrium constant expression (K) for the dissociation of ANY weak base can be written as:

�' = �� (�

�(

As this equilibrium constant describes the extent of dissociation of a weak acid in water, it’s known as the Base

Ionization Constant, Kb. The Kb values of several weak bases are shown below.


The pH of Commonly-Occurring Acid and Base Substances

Many common substances are either acids or bases. Citrus fruits such as oranges and lemons contain acid

compounds.

Vinegar contains acetic acid, HC2H3O2. This molecular compound is classified as a weak acid because when

dissolved in water, only a small fraction of the acetic acid molecules dissociate and donate their protons to water

molecules to make hydronium ions and the conjugate base of acetic acid, the acetate polyatomic ion, C2H3O2 -.

HC2H3O2 (aq) + H2O (l) � H3O+

(aq) + C2H3O2 -

(aq)

Bases can be found in antacid preparations and in many household cleaning solutions. The schematic below

indicates the pH of a number of familiar aqueous solutions.

Realize that if you know the pH of a solution, you can calculate the hydronium ion concentration from the

definition of pH:

If pH = – log [H3O+], then [H3O

+] = 10

– pH

(To use your calculator to find a pH, simply enter the hydronium ion concentration, [H3O+], then press the “log”

button, and then the +/- or (-) button. To do the reverse, that is, find a hydronium ion concentration from a pH,

simply enter the pH value, press the +/- or (-) button, and then the “anti-log” or “10x” button.)

Thus if the pH of blood is normally between 7.35 and 7.45, then the hydronium ion concentration in blood is

normally between 3.55 x 10-8

and 4.47 x 10-8

M (moles H3O+/liter blood solution):

If pH = 7.35 = – log [H3O+],

then [H3O+] = 10

– pH = 10

– 7.35 = 4.47 x 10

−8 M

And if pH = 7.45 = – log [H3O+],

then [H3O+] = 10

– pH = 10

– 7.45 = 3.55 x 10

−8 M


The Buffering Capability of Weak Acids and their Conjugate Bases

Solutions of weak acids and their conjugate bases are extremely useful in everyday life because of their ability to

resist changes in pH. Another way to say this is that they have “buffering” capability, or that they can serve as

“buffers” of the pH of an aqueous solution.

A buffer is a solution that contains two solutes: a weak acid, and the conjugate base of that weak acid.

These solutes are present in equal (or approximately equal) concentrations. (Remember that the conjugate base is

the polyatomic anion that is produced when a weak acid dissociates in water.)

For example, a buffer could be made of acetic acid, HC2H3O2, and its conjugate base, the acetate ion, C2H3O2 -.

HC2H3O2 (aq) + H2O (l) � H3O+

(aq) + C2H3O2 - (aq)

If we attempt to change the pH of a buffered solution by adding a strong acid (to lower the pH) or by adding a

strong base (to raise the pH), the buffer is able to resist the pH change. This is because the buffer contains a weak

base that is capable of reacting with any added acid, and it contains a weak acid that can react with any added base.

For example, if hydrochloric acid, HCl, is added to an acetic acid/acetate buffer, the C2H3O2 - ion (as the conjugate

base) can react with the HCl, preventing it from increasing the hydronium ion concentration and lowering the pH

of the solution:

C2H3O2

- (aq) + HCl (aq) → HC2H3O2 (aq) + Cl

- (aq)

Some of the C2H3O2 - ion is converted to HC2H3O2 by this reaction; thus, the concentration of the C2H3O2

- ion will

decrease with the addition of HCl, and the concentration of acetic acid will increase.

But note that the pH of the buffer solution will remain relatively constant (since the concentrations of H3O+ and

OH–

remain relatively constant).

The impact of an added strong base can also be minimized by this buffer solution. For example, sodium

hydroxide, NaOH, will react with the HC2H3O2 in the buffer (converting it to C2H3O2 - ion), rather than increasing

the hydroxide ion concentration in the solution and raising the pH.

HC2H3O2 (aq) + NaOH (aq) → Na+ (aq) + C2H3O2

- (aq) + H2O (aq)

The above reactions will occur and maintain the pH of the solution as long as there are enough acetate ions

(C2H3O2-) to react with any added acid, or enough acetic acid molecules (HC2H3O2) to react with any added base.

If either the weak acid or conjugate base is completely used up in these reactions, then the buffer’s “capacity” has

been exceeded, and any additional acid or base will cause the pH to change.

Thus the higher the concentrations of the weak acid and conjugate base in the buffer system, the greater the buffer

capacity, and the more moles of strong acid or base that can be added to the buffer without causing a significant

change in pH.

Buffers are critically important in biological situations. The human body uses a variety of buffering systems to

maintain both extracellular and intracellular solutions at the optimal pH. For example, the carbonic acid –

bicarbonate ion buffering system (H2CO3 – HCO3-) is critical to maintaining the pH of the blood between 7.35 and

7.45.


Indicators of pH

An acid-base indicator is a molecular compound that undergoes a color change when the pH of the solution

containing the indicator reaches a certain value. Some common indicator compounds are listed in the figure

below. You’ll note that the different indicators change color at different pH ranges; this allows chemists to select a

specific indicator that will change color at precisely the desired pH.

Indicator solutions can be added directly to solutions of acids or bases. Alternatively, strips of filter paper can be

impregnated with indicator compounds to produce “pH paper.” When a drop of acid or base solution is touched to

the pH paper, the paper will change colors. By matching the color of the paper to an indicator color chart, the pH

of the solution can be determined.

In today’s experiment, you’ll use an acid-base indicating paper to:

• Determine the acidity or basicity of some common household substances

• Compare the pH of two different concentrations of strong acid and strong base solutions

• Compare the pH of strong and weak acid solutions that have equal concentrations

• Compare the pH of strong and weak base solutions that have equal concentrations

• Investigate the impact of adding a strong acid or strong base to a buffer system.

pHydrion Acid-Base Indicating Paper

The acid-base indicating paper you’ll use in this

experiment pH impregnated with a “universal indicator”

solution that is able to turn one of fourteen different

colors in response to pH values from 0 (extremely

acidic) to 13 (extremely basic). As the pH of tested

solutions vary from acidic to basic, the response colors

of the pHydrion range from red (pH = 0) to orange to

yellow to light green (pH = 7) to dark green to light blue

to dark blue (pH = 13). The various hues are too

complex to describe here, so the test strips of the pH

paper must be compared with a pHydrion color chart.


Procedure that you will follow

A. pH of Common Household Substances

1. Examine the clear plastic well plate in your laboratory equipment drawer. It should be of one of two designs:

• Three rows (labeled A, B, C) containing four wells (labeled 1, 2, 3, 4) in each row, OR

• Four rows (labeled A, B, C, D) containing six wells (labeled 1, 2,3, 4, 5, 6) in each row.

2. You’ll obtain a sample of each of the nine common household substances listed below (and in the Data/Results

table in the Report Sheet) in a separate well in the well plate. To keep from confusing the samples with one

another, as you add each sample to a well, record the label of that well (ex: A-1 for row A, well #1) on the

Report Sheet.

The substances are:

• Carbonated Soda

• White Vinegar

• Lemon Juice

• Distilled Water

• Sodium Chloride

• Table Sugar (Sucrose)

• Sodium Bicarbonate

• Antacid Tablet

• Household Ammonia

3. Obtain the samples as follows:

• For liquids: Use a pipet or eye-dropper to add 10 drops of the solution to the well.

• For solids: Use a spatula to place a sample about the size of a match head into the well (some solids may

need to be crushed). Add 10 drops of distilled water to the sample in the well and stir to dissolve.

4. Test for the actual pH value of each sample with pHydrion paper by touching a glass stirring rod to the

solution, then touching the rod to the paper (be sure to wipe the rod clean between samples). In this way, the

same piece of pHydrion paper can be used to test several different samples. Use the color chart provided with

the pHydrion paper to determine the actual pH. Note: The color change of the pH paper must be observed

while the paper is still wet with the tested solution.

5. Dispose of the samples by rinsing the contents of the well plate into a large waste beaker, then pouring this into

the appropriate waste container.

B. Determining the pH of Strong and Weak Acids and Bases

1. Place a 1 mL sample of each the following acid or base solutions in a separate well in the well plate. Again, to

keep from confusing the solutions with one another, as you place each sample in a well, record the label of that

well (ex: A-1 for row A, well #1) on the Report Sheet.

The solutions are:

• 0.10 M HCl (hydrochloric acid, a strong acid)

• 0.00010 M HCl (hydrochloric acid, a strong acid)

• 0.10 M NaOH (sodium hydroxide, a strong base)

• 0.00010 M NaOH (sodium hydroxide, a strong base)

• 0.10 M HC2H3O2 (acetic acid, a weak acid)

• 0.10 M NaC2H3O2 (sodium acetate, a weak base)


• 0.10 M NH3 (ammonia, a weak base)

• 0.10 M NH4Cl (ammonium chloride, a weak acid)

2. Test for the actual pH value of each solution with pHydrion paper. As before, use the glass stirring rod to test

several different samples on the same piece of pH paper. Use the color chart provided with the container of

pHydrion paper to determine the actual pH.



C. Determining the Effect of a Buffer

1. Place the samples listed below into separate wells in the well plate as indicated by the labels below.

• Buffer in well A-1 (Acetic Acid/Acetate): 1 mL of 0.10 M HC2H3O2 and 1 mL of 0.10 M NaC2H3O2

• Buffer in well A-2 (Acetic Acid/Acetate): 1 mL of 0.10 M HC2H3O2 and 1 mL of 0.10 M NaC2H3O2

• Buffer in well B-1 (Ammonium/Ammonia): 1 mL of 0.10 M NH4Cl and 1 mL of 0.10 M NH3

• Buffer in well B-2 (Ammonium/Ammonia): 1 mL of 0.10 M NH4Cl and 1 mL of 0.10 M NH3

• Control in well C-1: 2 mL of distilled water

• Control in well C-2: 2 mL of distilled water

2. Test for the actual pH value of each buffer solution (and the control samples) with pHydrion paper. As before,

use the glass stirring rod to test several different samples on the same piece of pH paper. Use the color chart

provided with the container of pHydrion paper to determine the actual pH.

3. To the samples in wells A-1, B-1, and C-1: Add 1 drop of the 0.10 M HCl solution. Mix the samples carefully

by gently swirling the well plate. Test for the actual pH of each sample with pHydrion paper (using the glass

stirring rod to test the three samples on the same piece of pH paper, and comparing the colors with the chart to

determine the actual pH).

4. To the samples in wells A-2, B-2, and C-2: Add 1 drop of the 0.10 M NaOH solution. Mix the samples

carefully by gently swirling the well plate. Test for the actual pH of each sample with pHydrion paper (using

the glass stirring rod to test the three samples on the same piece of pH paper, and comparing the colors with the

chart to determine the actual pH).




Report Sheet 11: Introduction to Acids, Bases, pH, and Buffers Student ______________________________ Lab Partner__________________________ Date Lab Performed__________

Section #_________ Lab Instructor__________________________________________ Date Report Received ___________

Lab Notebook: Data and Observations

A. Determining the pH of Common Substances

For each one of the common substances tested:

Were the pH results you obtained (relative acidity or basicity) consistent with what you would expect for that

substance?

Did you expect the substance to be an acid? A base? Did you expect the pH to be that high? That low?

Explain how the results compared to your expectations for each substance

Substance

Spot

Plate

Code

Color

pHydrion

Paper

pHydrion

pH

Explain how results

compare to expectations

Carbonated Soda

White Vinegar

Lemon Juice

Distilled water

Sodium Chloride

Sugar (sucrose)

Sodium Bicarbonate

Antacid Tablet

Household Ammonia


B. Determining the pH of Strong and Weak Acids and Bases

Solution

Spot

Plate

Code

Known to be acid

or base?

Known to be weak

or strong?

Color pHydrion

Paper pHydrion pH

0.10 M HCl

0.00010 M HCl

0.10 M NaOH

0.00010 M NaOH

0.10 M HC2H3O2

0.10 M NaC2H3O2

0.10 M NH3

0.10 M NH4Cl

C. Determining the Effect of a Buffer

BEFORE addition

of HCl or NaOH

AFTER addition

of HCl or NaOH

Buffer Solution

Spot

Plate

Code

Color

pHydrion

Paper

pHydrion

pH

Color

pHydrion

Paper

pHydrion

pH Change in pH

pHfinal – pHinitial

0.10 M HC2H3O2

+

0.10 M NaC2H3O2

A-1

0.10 M HC2H3O2

+

0.10 M NaC2H3O2

A-2

0.10 M NH4Cl +

0.10 M NH3

B-1

0.10 M NH4Cl +

0.10 M NH3

B-2

Distilled water C-1

Distilled water C-2


Formal Report: Results and Conclusions for Effect of Concentration

1. Compare the pH of the two different concentrations of the strong acid HCl:

What is the [H3O+] in the 0.10 M HCl

solution?

And what should theoretically be the pH of the

0.10 M HCl solution? Show your work.

What was the

measured pH of the

0.10 M HCl solution?

What is the [H3O+] in the 0.00010 M

HCl solution?


0.00010 M HCl solution? Show your work.

What was the

measured pH of the

0.00010 M HCl

solution?

2. Compare the pH of the two different concentrations of the strong base NaOH:


NaOH solution? Show your work.


0.10 M NaOH solution? Show your work.

What was the

measured pH of the

0.10 M NaOH

solution?


NaOH solution? Show your work.


0.00010 M NaOH solution? Show your work.

What was the

measured pH of the

0.00010 M NaOH

solution?

3. What relationship did you observe between the concentration of the strong acid and the pH of that solution?

Between the concentration of the strong base and the pH of that solution?


4. List and then compare the measured pH values of the three 0.10 M acid solutions:

• hydrochloric acid (HCl) _____________________

• acetic acid (HC2H3O2) _____________________

• ammonium chloride (NH4Cl) _____________________

These three solutions have the SAME concentration, but do they have the SAME pH? _____________________

You already know that HCl is a strong acid. Of the remaining two acids (both of which are weak acids):

Which weak acid has the higher concentration of hydronium ions? ________________________________

Which weak acid is stronger than the other? ________________________________

Which weak acid would have the larger Ka value (acid dissociation constant)? ____________________________

Briefly state how you reached these conclusions.

5. List and then compare the measured pH values of the three 0.10 M base solutions:

• sodium hydroxide (NaOH) _____________________

• sodium acetate (NaC2H3O2) _____________________

• ammonia (NH3) _____________________

These three solutions have the SAME concentration, but do they have the SAME pH? _____________________

You already know that NaOH is a strong base. Of the remaining two bases (both of which are weak bases):

Which weak base has the higher concentration of hydroxide ions? ________________________________

Which weak base is stronger than the other? ________________________________

Which weak base would have the larger Kb value (base ionization constant)? _____________________________

Briefly state how you reached these conclusions.

6. List and then compare the change in pH when HCl or NaOH was added to each solution below:

• acetic acid (HC2H3O2) + sodium acetate (NaC2H3O2): ∆pH with HCl =________ ∆pH with NaOH =_________

• ammonium chloride (NH4Cl) + ammonia (NH3): ∆pH with HCl =_________ ∆pH with NaOH =_________

• distilled water: ∆pH with HCl =_________ ∆pH with NaOH =_________

Are these results appropriate for the addition of HCl/NaOH to a buffer solution (vs. pure water)? Explain.

11. introduction to acids, bases, ph, and · pdf fileintroduction to acids, bases, ph, ... the...

Documents