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CHEM 404
Exp 7
Experiment 7:
pH MEASUREMENTS – BUFFERS AND THEIR PROPERTIES*
One of the more important properties of an aqueous solution is its concentration of
hydrogen ion. The H+ or H3O+ ion has great effect on the solubility of many inorganic and organic
species, on the nature of complex metallic cations found in solutions, and on the rates of many
chemical reactions. It is important to know how to measure the concentration of hydrogen ion and
understand its effect on solution properties.
For convenience, the concentration of H+ ion is frequently expressed as the pH of the
solution rather than as molarity. The pH of a solution is defined by the following equation:
pH = -log[H+]
(1)
where the logarithm is taken to the base 10. If [H+] is 1 x 10-4 moles per liter, the pH of the solution
is, by the equation, 4. If [H+] is 5 x 10-2 M, the pH is 1.3.
Basic solutions can also be described in terms of pH. In water solutions, the following
equilibrium relation will always be obeyed:
Kw = [H+][OH-] = 1 x 10-14 at 25 °C
(2)
In distilled water, [H+] equals [OH-], so, by Equation 2, [H+] must be 1 x 10-7 M. Therefore, the
pH of distilled water is 7. Solutions in which [H+] is greater than [OH-] are said to be acidic and
will have a pH less than 7; if [H+] is less than [OH-], the solution is basic and will have a pH greater
than 7. Combining Equations 1 and 2, a solution with a measured pH of 10 will have a [H+] of 1 x
10-10 M and a [OH-] of 1 x 10-4 M.
The pH of a solution can be experimentally measured in two ways. In the first of these, a
chemical called an indicator, which is sensitive to pH, is used. These substances have colors that
change over a relatively short pH range (about two pH units) and can, when properly chosen, be
used to determine roughly the pH of a solution. Two very common indicators are litmus, usually
used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus
changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein
changes from colorless to red as the pH goes from 8 to 10. A given indicator is useful for
determining pH only in the region in which it changes color. Indicators are available for
measurement of pH in all the important ranges of acidity and basicity. By matching the color of a
suitable indicator in a solution of known pH with that in an unknown solution, one can determine
the pH of the unknown to within about 0.3 pH units.
The other method for finding pH is with a device called a pH meter. In this device two
electrodes, one of which is sensitive to [H+], are immersed in a solution. The potential between the
two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish
the pH of the solution. A pH meter gives much more precise measurement of pH than does a typical
indicator and is ordinarily used when an accurate determination of pH is needed.
*Adapted from Slowinski, E. J., Wolsey, W. C. Chemical Principles in the Laboratory 9th ed.
CHEM 404
Exp 7
Some acids and bases undergo substantial ionization in water, and are called strong because
they essentially completely ionize in reasonably dilute solutions. Other acids and bases, because
of incomplete ionization (often only about 1% in 0.1 M solution), are called weak. Hydrochloric
acid, HCl, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base.
Acetic acid, HC2H3O2, and ammonia, NH3, are classic examples of a weak acid and a weak base.
A weak acid will ionize according to the Law of Chemical Equilibrium:
HA(aq) ? H+(aq) + A-(aq)
At equilibrium,
Ka =
[H+ ][A- ]
[HA]
(3)
(4)
Ka is a constant characteristic of the acid HA; in solutions containing HA, the product of
concentrations in the equation will remain constant at equilibrium independent of the manner in
which the solution was made. A similar relation can be written for solutions of a weak base.
The value of the ionization constant Ka for a weak acid can be found experimentally in
several ways. In general, the concentrations of each of the species in Equation 4 need to be found
by one means or another. In this experiment, you will determine Ka for a weak acid in connection
with your study of the properties of solutions called buffers.
Salts that can be formed by the reaction of strong acids and strong bases – such as NaCl,
KBr, or NaNO3 – ionize completely but do not react with water when in solution. They form
neutral solutions with a pH of about 7. When dissolved in water, salts of weak acids or weak bases
furnish ions that tend to react to some extent with water, producing molecules of the weak acid or
base and liberating some OH- or H+ ion to the solution.
If HA is a weak acid, the A- ion produced when NaA is dissolved in water will react with
water to some extent, according to the equation:
A-(aq) + H2O ? HA(aq) + OH-(aq)
(5)
Solutions of sodium acetate, NaC2H3O2, the salt formed by reaction of sodium hydroxide with
acetic acid, will be slightly basic because of the reaction of ion with water to produce HC 2H3O2
and OH-. Because of the analogous reaction of the ion with water to form H3O+ ion, solutions of
ammonium chloride, NH4Cl, will be slightly acidic.
Salts of most transition metal ions are acidic. A solution of CuSO4 or FeCl3 will typically
have a pH equal to 5 or lower. The salts are completely ionized in solution. The acidity comes
from the fact that the cation is hydrated (e.g. Cu(H2O)42+, or Fe(H2O)63+). The large + charge on
the metal cation attracts electrons from the O–H bonds in water, weakening them and producing
some H+ ions in solution; with CuSO4 solutions the reaction would be:
Cu(H2O)42+(aq) ? Cu(H2O)3OH+(aq) + H+(aq)
(6)
CHEM 404
Exp 7
Some solutions, called buffers, are remarkably resistant to pH changes. Water is not a
buffer, since its pH is very sensitive to addition of any acidic or basic species. Even bubbling your
breath through a straw into distilled water can lower its pH by at least one pH unit, just due to the
small amount of CO2, which is acidic, in exhaled air. With a good buffer solution, you could blow
exhaled air into it for half an hour and not change the pH appreciably. All living systems contain
buffer solutions, since stability of pH is essential for the occurrence of many of the biochemical
reactions that go on to maintain the living organism.
There is nothing mysterious about what one needs to make a buffer. All that is required is
a solution containing a weak acid and its conjugate base. An example of such a solution is one
containing the weak acid HA, and its conjugate base, A- ion, obtained by dissolving the salt NaA
in water.
The pH of such a buffer is established by the relative concentrations of HA and A- in the
solution. By rearranging Equation 4, you can solve for the concentration of H+ ion:
[H+] = Ka x
[HA]
[A- ]
(4a)
This equation is often further modified, by taking the negative logarithm to base 10 of both sides:
pH = pKa + log
[A- ]
[HA]
(4b)
where pKa equals the negative log of Ka. This equation is called the Henderson-Hasselbach
equation, and is set up to conveniently calculate the resulting pH of a buffer from know quantities
of a weak acid and its conjugate base.
If you are working with a weak acid HA whose Ka equals 1 x 10-5, its pKa equals 5.0. If
you mix equal volumes of 0.10 M HA and 0.10 M NaA, the pH equals pKa equals 5.0, since the
concentrations of HA and A- are equal in the final solution, and the log of 1 equals 0.
You might wonder why [HA] and [A-] do not change when the species are mixed. Actually,
they do, but only very slightly, just enough to generate enough H+ ion to satisfy the condition
imposed by Equation 4a. Ordinarily, Ka is very small, so the amount of H+ produced from
ionization of HA in solution is also very small. Since only a negligible decrease in [HA] occurs,
and only a tiny increase in [A-] results, you can work from the assumption that in any buffer the
weak acid and its conjugate base do not react appreciably with one another when their solutions
are mixed. Their relative concentrations can be calculated from the way the buffer was put together.
Using Equation 4b, you should be able to answer several questions regarding buffers. For
example, how does the pH of buffer solution change if you dilute it with water? How does it change
if you add more HA solution? What happens if you add more solution of NaA, containing the
conjugate base? What is the pH range over which a buffer would be useful, if you assume that the
ratio of [A-] to [HA] must lie between 10:1 and 1:10? How sensitive is the pH of water to the
addition of a strong acid, like HCl, or addition of a strong base, like NaOH? What happens when
CHEM 404
Exp 7
you add a strong acid or a strong base to a buffer? Why does the pH of a buffer resist changing
when small amounts of a strong acid or strong base are added?
Some of these questions you can answer by just looking at Equation 4b. For others, you
need to consider that, although a weak acid and its conjugate base don’t react with each other, a
weak acid like HA will react quantitatively in a buffer with a strong base, like NaOH:
HA(aq) + OH-(aq) ? A-(aq) + H2O
(7)
and a weak base like NaA will react with a strong acid like HCl:
A-(aq) + H+(aq) ? HA(aq)
(8)
In Reaction 7, a small amount of NaOH added will be “soaked up” by the acid HA, producing
some B- ion and increasing the value of [A-]/[HA] and the pH slightly, but not destroying the
buffer. In Reaction 8, a similar interaction between added HCl and A- will decrease the value of
[A-]/[HA] and the pH slightly.
If you continued to add HCl to the buffer, eventually, through Reaction 8, the strong acid
would react with all of the A- present, and the buffer would be “exhausted” since it would contain
only HA. At that point, any excess HCl would produce a pH with just about the same value as if
the HCl were added to water. Similar behavior would occur if you added a large amount of NaOH
to the buffer. Typically, the range over which a buffer is useful is limited to about +/- one pH unit
from the pKa of the weak acid component.
In this experiment, you will determine the approximate pH of several solutions by using
acid-base indicators. Then you will find the pH of some other solutions with a pH meter. In the
rest of the experiment you will carry out some reactions that will allow you to answer all the
questions raised about buffers. Finally, you will prepare one or two buffers having specific pH
values.
Procedure
For Parts A and B, you may work in groups of up to four at your lab bench to collect data
and observations for the solutions being studied. For Part C, you will work with your lab partner.
A: Determination of pH by Use of Acid-Base Indicators
To each of five small test tubes add about 1 mL 0.10 M HCl. To each tube add a drop or
two of one of the indicators listed in Table 1, one indicator to a tube. Note the color of the solution
you obtain in each case. By comparing the colors you observe with the information in Table 1,
estimate the pH of the solution to within a range of one pH unit, for example “1-2”, or “4-5”. In
making your estimate, note that the color of an indicator is most indicative of pH in the region
where the indicator is changing color.
Rinse your test tubes well with distilled water, then repeat this procedure with each of the
following solutions:
0.10 M NaH2PO4
0.10 M HC2H3O2
0.10 M ZnSO4
CHEM 404
Exp 7
Table 1: pH Range of Some Chemical Indicators
B. Measurement of the pH of Some Typical Solutions
For the rest of the experiment, you will be using a pH meter to measure the pH of your
solutions. Your instructor will show you how to properly operate the pH meter. Note that the
electrodes are fragile, so use care when handling the electrode probe.
Obtain 25 mL of each of the following 0.1 M solutions in labeled 150-mL beakers:
NaCl
Na2CO3
NaC2H3O2
NaHSO4
Measure and record the pH of each solution, taking care to rinse the electrode with distilled water
between measurements. Some of these solutions will be nearly neutral, others significantly acidic
or basic. For each solution having a pH less than 6 or greater than 8, write a net ionic equation that
explains qualitatively why the observed pH value is reasonable.
After you have completed your pH measurements, add a drop or two of Bromcresol green
to your solutions and record the color you observe. Write a rationale for the colors obtained with
Bromcresol green for these solutions.
C. Some Properties of Buffers
You will be provided with 0.10 M stock solutions that can be used to make three different
common buffer systems:
HC2H3O2–C2H3O2acetic acid–acetate ion
NH4+–NH3
ammonium ion–ammonia
HCO3—CO32hydrogen carbonate–carbonate
The sources of the ions will be sodium and ammonium salts containing those ions. Select one of
these buffer systems for your experiment. In a clean, dry, labeled beaker, obtain approximately 40
mL of your acid component. The acid will be one of the following solutions: 0.10 M HC2H3O2,
0.10 M NH4Cl, or 0.10 M NaHCO3.
In a second clean, dry, labeled beaker, obtain approximately 60 mL of your base component.
The base will be one of the following solutions: 0.10 M NaC2H3O2, 0.10 M NH3, or 0.10 M Na2CO3.
CHEM 404
Exp 7
In a 25-mL graduated cylinder, measure 15 mL of your acid component and add it to
another clean, dry beaker. Rinse the graduated cylinder with distilled water, then measure 15 mL
of your base component and add it to the 15 mL of acid.
Stir the buffer solution that you have just prepared, measure the pH and record it on the
Data page. Calculate the pKa for the acid. Add 30 mL of water to your buffer mixture, stir, and
pour half of the resulting solution into another beaker. Measure the pH of the diluted buffer.
Calculate the pKa once again.
Add five drops of 0.10 M HCl to one of your diluted buffer solutions, then measure the
resulting pH. To the other beaker of the diluted buffer add 5 drops 0.10 M NaOH, and again
measure the pH. Record your results.
Measure 15 mL distilled water into a 100-mL beaker. Measure the pH. Add 5 drops 0.10
M HCl and measure the pH again. To that solution add 10 drops 0.10 M NaOH, mix, and measure
the pH. Compare the pH changes between your buffer system and water.
Prepare a buffer mixture containing 2 mL of the acid component and 20 mL of the base
component. Mix, and measure the pH. Calculate a third value for pKa. To that solution add 3 mL
0.10 M NaOH, which should exhaust the acid component of your buffer. Measure the pH. Explain
your results.
Select a pH different from any of those you observed in your experiments. Using Equation
4b, calculate the appropriate ratio of your basic and acidic component required to prepare that
buffer. Obtain and mix the appropriate volumes of your acidic and basic components to make up
the buffer, then measure its pH.
CHEM 404 – Exp 7
Name:______________________________ Section: _________
Data & Calculations
A. Determination of pH by Use of Acid-Base Indicators
Indicator
Color with 0.1 M Solutions
HCl
NaH2PO4
HC2H3O2
ZnSO4
Methyl violet
_____________ _____________ _____________ _____________
Thymol blue
_____________ _____________ _____________ _____________
Methyl yellow
_____________ _____________ _____________ _____________
Congo red
_____________ _____________ _____________ _____________
Bromcresol green
_____________ _____________ _____________ _____________
pH range
_____________ _____________ _____________ _____________
Circle the color observation for each solution that was most useful in estimating the pH range.
B. Measurement of the pH of Some Typical Solutions
Record the pH and the color observed with Bromcresol green for each of the 0.1 M solutions that
were tested.
NaCl
Na2CO3
NaC2H3O2
NaHSO4
pH
_____________
_____________
_____________
_____________
color
_____________
_____________
_____________
_____________
For any solutions with a pH lower than 6 or higher than 8, write a net ionic equation, similar to
Reaction 5, to explain qualitatively why the solution has that pH.
Solution _________ Equation _____________________________________________________
Solution _________ Equation _____________________________________________________
Explain why the color observed with Bromcresol green for each of the four solutions is reasonable,
given the pH.
C. Some Properties of Buffers
Buffer system selected: Acid ________________
Conjugate Base ________________
pH of prepared buffer:
____________
pKa: ____________
Ka: ____________
pH of diluted buffer:
____________
pKa: ____________
Ka: ____________
pH after 5 drops HCl:
____________
pH after 5 drops NaOH:
____________
pH of distilled water:
____________
pH after 5 drops HCl:
____________
pH after 10 drops NaOH:
____________
pKa: ____________
Ka: ____________
pH of buffer in which
pH after 3 mL NaOH:
[A- ]
[HA]
= 10: ____________
____________
Explain your observations after 3 mL of NaOH was added to this buffer. What occurred in the
solution to give the pH change you observed?
Average value from all buffer solutions:
pKa: ____________
Ka: ____________
Target pH of buffer to prepare: ____________
Ratio of
[A- ]
[HA]
required (from Equation 4b, and average pKa): ____________
Volume 0.1 M NaA used: ____________ mL
Measured pH of prepared buffer: ____________
Volume of 0.1 M HA used: ____________ mL
CHEM 404 – Exp 7
Name:______________________________ Section: _________
Prelab Assignment
1. A solution of a weak acid was tested with the indicators used in this experiment. The colors
observed were as follows:
Methyl violet: violet
Thymol blue: orange
Methyl yellow: red
Congo red: violet
Bromcresol green: yellow
What is the approximate pH of the solution? ____________
2. An aqueous solution of NH3 has a pH of 11.6. The ammonia (NH3) molecule is the conjugate
base of the ammonium (NH4+) ion. Write the net ionic equation for the reaction between ammonia
and water that makes an aqueous solution of NH3 basic. (Similar to Reaction 5.)
3. The pH of a 0.10 M HCN solution is 5.2.
a. What is [H+] in the solution?
____________ M
b. What is [CN-]? What is [HCN]? (Where do the H+ and CN- ions come from in the solution?)
[CN-] = ____________ M; [HCN] = ____________ M
c. What is the value of Ka for HCN (use Equation 4)? What is the pKa for HCN?
Ka = ____________; pKa = ____________
4. Formic acid, HCO2H, has a Ka value of 1.8 x 10-4. A student is asked to prepare a buffer with a
pH of 3.90 from a 0.10 M solution of formic acid and a 0.10 M solution of sodium formate. How
many milliliters of the sodium formate solution should she add to 20. mL of the formic acid
solution to make the buffer? (use Equation 4a or 4b)
____________ mL

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