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Principle; Buffers*
Cr2O7 2–(aq) + H2O(l)
Jo A. Beran/Trey Hernandez
2 CrO4 2–(aq) + 2H+(aq)
The chromate ion (left) is yellow, and the dichromate ion (right) is
orange. An equilibrium between the two ions is affected by changes
in pH.
• To study the effects of concentration and temperature changes on the position of
equilibrium in a chemical system
• To study the effect of strong acid and strong base addition on the pH of buffered
and unbuffered systems
• To observe the common-ion effect on a dynamic equilibrium
The following techniques are used in the Experimental Procedure:
Most chemical reactions do not produce a 100% yield of product, not because of
experimental technique or design, but rather because of the chemical characteristics
of the reaction. The reactants initially produce the expected products, but after a
period of time the concentrations of the reactants and products stop changing.
This apparent cessation of the reaction before a 100% yield is obtained implies
that the chemical system has reached a state where the reactants combine to form the
products at a rate equal to that of the products re-forming the reactants. This condition
is a state of dynamic equilibrium and is characteristic of all reversible reactions.
For the reaction
2 NO2(g)
N2O4(g) + 58 kJ
chemical equilibrium is established when the rate at which two NO2 molecules react
equals the rate at which one N2O4 molecule dissociates (Figure 16.1).
If the concentration of one of the species in the equilibrium system changes, or if
the temperature changes, the equilibrium tends to shift in a way that compensates for the
change. For example, assuming the system represented by equation 16.1 is in a state of
dynamic equilibrium, if more NO2 is added, the probability of its reaction with other
NO2 molecules increases. As a result, more N2O4 forms, and the reaction shifts to the
right, until equilibrium is reestablished.
A general statement governing all systems in a state of dynamic equilibrium follows:
Figure 16.1 A dynamic
equilibrium exists between
reactant NO2 molecules and
product N2O4 molecules.
If an external stress (change in concentration, temperature, etc.) is
applied to a system in a state of dynamic equilibrium, the equilibrium
shifts in the direction that minimizes the effect of that stress.
*Numerous online Web sites discuss LeChâtelier’s principle.
Experiment 16
This is LeChâtelier’s principle, proposed by Henri Louis LeChâtelier in 1888.
Often the equilibrium concentrations of all species in the system can be
determined. From this information, an equilibrium constant can be calculated;
its magnitude indicates the relative position of the equilibrium. This constant is
determined in Experiments 22, 26, and 34 .
Two factors affecting equilibrium position are studied in this experiment: changes
in concentration and changes in temperature.
Changes in Concentration
Complex ion: a metal ion bonded to
a number of Lewis bases. The
complex ion is generally identified by
its enclosure with brackets, [ ].
Metal–Ammonia Ions. Aqueous solutions of copper ions and nickel ions appear sky
blue and green, respectively. The colors of the solutions change, however, in the presence of added ammonia, NH3. Because the metal–ammonia bond is stronger than the
metal–water bond, ammonia substitution occurs and the following equilibria shift
right, forming the metal–ammonia complex ions:1
[Cu(H2O)4]2+(aq) + 4 NH3(aq)
[Cu(NH3)4]2+(aq) + 4 H2O(l)
[Ni(H2O)6]2+(aq) + 6 NH3(aq)
[Ni(NH3)6]2+(aq) + 6 H2O(l)
Addition of strong acid, H , affects these equilibria by its reaction with ammonia
(a base) on the left side of the equations:
NH3(aq) + H+(aq) —› NH4 +(aq)
Michael Watson
The ammonia being removed from the equilibria causes the reactions to shift left
to relieve the stress caused by the removal of the ammonia, re-forming the aqueous
Cu2+ (sky blue) and Ni2+ (green) solutions. For copper ions, this equilibrium shift may
be represented as
[Cu(H2O)4]2+(aq) + 4 NH3(aq)
[Cu(NH3)4]2+(aq) + 4 H2O(l)
4 H (aq)
4 NH4 +(aq )
Multiple Equilibria with
the Silver Ion
[Cu(H2O)4]2+ is a sky-blue color
(left), but [Cu(NH3)4]2+ is a deepblue color (right).
Many salts are only slightly soluble in water. Silver ion, Ag+, forms a number of these
salts. Several equilibria involving the relative solubilities of the silver salts of the carbonate, CO3 2–, chloride, Cl–, iodide, I–, and sulfide, S2–, anions are investigated in this
Silver Carbonate Equilibrium. The first of the silver salt equilibria observed in this
experiment is that of a saturated solution of silver carbonate, Ag2CO3, in dynamic
equilibrium with its silver and carbonate ions in solution.
2 Ag+(aq) + CO3 2–(aq)
Nitric acid, HNO3, dissolves silver carbonate: H+ ions react with (and remove) the
CO3 2– ions on the right; the system, in trying to replace the CO3 2– ions, shifts to
the right. The Ag2CO3 dissolves, and carbonic acid, H2CO3, forms.
2 Ag+(aq) + CO3 2–(aq)
2 H+(aq )
H2CO3(aq ) —› H2O(l) + CO2(g)
The carbonic acid, being unstable at room temperature and pressure, decomposes
to water and carbon dioxide. The silver ion and nitrate ion (from HNO3) remain in
A further explanation of complex ions appears in Experiment 36.
LeChâtelier’s Principle; Buffers
Silver Chloride Equilibrium. Chloride ion precipitates silver ion as AgCl. Addition of
chloride ion (from HCl) to the above solution containing Ag+ causes the formation of
a silver chloride, AgCl, precipitate, now in dynamic equilibrium with its Ag+ and Cl–
ions (Figure 16.2).
Ag+(aq) + Cl–(aq)
Aqueous ammonia, NH3, “ties up” (i.e., it forms a complex ion with) silver ion,
producing the soluble diamminesilver(I) ion, [Ag(NH3)2]+. The addition of NH3
removes silver ion from the equilibrium in equation 16.9, shifting its equilibrium position to the left and causing AgCl to dissolve:
Adding acid, H+, to the solution again frees silver ion to recombine with chloride ion
and re-forms solid silver chloride. This occurs because H+ reacts with the NH3 (see equation 16.4) in equation 16.10, restoring the presence of free Ag+ to combine with the free
Cl– to form AgCl(s) shown in equation 16.9.
Figure 16.2 Solid AgCl quickly
forms when solutions containing
Ag+ and Cl– are mixed.
Ag+(aq ) + Cl–(aq)
2 NH3(aq) + 2 H+(aq ) —› 2 NH4 +(aq)
Ken Karp
Ag+(aq) + Cl–(aq)
2 NH3(aq )
[Ag(NH3)2]+(aq )
Silver Iodide Equilibrium. Iodide ion, I– (from KI), added to the Ag+(aq) + 2 NH3(aq)
Ag(NH3)2 + (aq) equilibrium in equation 16.10 results in the formation of solid
silver iodide, AgI.
Ag+(aq) + 2 NH3(aq)
I –(aq )
The iodide ion removes the silver ion, causing a dissociation of the [Ag(NH3)2]+
ion and a shift of the equilibrium to the left.
Silver Sulfide Equilibrium. Silver sulfide, Ag2S, is less soluble than silver iodide, AgI.
Therefore, an addition of sulfide ion (from Na2S) to the AgI(s)
Ag+(aq) + I–(aq)
dynamic equilibrium in equation 16.12 removes silver ion; AgI dissolves, but solid
silver sulfide forms.
Ag+(aq) + I– (aq)
⁄2 S2–(aq )
⁄2 Ag2S(s)
In many areas of research, chemists need an aqueous solution that resists a pH change
when small amounts of acid or base are added. Biologists often grow cultures that are very
susceptible to changes in pH and therefore a buffered medium is required (Figure 16.3,
page 210).
A buffer solution must be able to consume small additions of H3O+ and OH– without undergoing large pH changes. Therefore, it must have present a basic component
that can react with added H3O+ and an acidic component that can react with added
OH–. Such a buffer solution consists of a weak acid and its conjugate base (or weak
base and its conjugate acid). This experiment shows that the acetic acid–acetate buffer
system can minimize large pH changes:
CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3CO2 –(aq)
Experiment 16
Center for Disease Control
The addition of OH– shifts the buffer equilibrium, according to LeChâtelier’s principle, to the right because of its reaction with H3O+, forming H2O. The shift right is by
an amount that is essentially equal to the moles of OH– added to the buffer system.
Thus, the amount of CH3CO2 – increases, and the amount of CH3COOH decreases by
an amount equal to the moles of OH– added:
CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3CO2 –(aq)
OH–(aq )
2 H2O(l)
H3O+(aq) + CH3CO2 –(aq)
H3O+(aq )
CH3COOH(aq) + H2O(l)
Figure 16.3 Bacteria cultures
survive in media that exist over a
narrow pH range. Buffers are
used to control large changes
in pH.
Conversely, the addition of H3O+ from a strong acid to the buffer system causes
the equilibrium to shift left, the H3O+ combines with the acetate ion (a base) to form
more acetic acid, an amount (moles) equal to the amount of H3O+ added to the system.
As a consequence of the addition of strong acid, the amount of CH3COOH
increases, and the amount of CH3CO2 – decreases by an amount equal to the moles of
strong acid added to the buffer system.
This experiment compares the pH changes of a buffered solution to those of an
unbuffered solution when varying amounts of strong acid or base are added to each.
Common-Ion Effect
The effect of adding an ion or ions common to those already present in a system at a
state of dynamic equilibrium is called the common-ion effect. The effect is observed
in this experiment for the following equilibrium:
4 Cl–(aq) + [Co(H2O)6]2+(aq)
[CoCl4]2–(aq) + 6 H2O(l)
Ligand: a Lewis base that donates a
lone pair of electrons to a metal ion,
generally a transition metal ion (see
Experiment 36).
Equation 16.17 represents an equilibrium of the ligands Cl– and H2O bonded to
the cobalt(II) ion—the equilibrium is shifted because of a change in the concentrations
of the chloride ion and water.
Changes in Temperature
Referring again to equation 16.1,
2 NO2(g)
Exothermic: characterized by energy
release from the system to the
Coordination sphere: all ligands of
the complex ion (collectively with the
metal ion they are enclosed in square
brackets when writing the formula of
the complex ion). See Experiment 36.
N2O4(g) + 58 kJ
(repeat of 16.1)
The reaction for the formation of colorless N2O4 is exothermic by 58 kJ. To favor
the formation of N2O4, the reaction vessel should be kept cool (Figure 16.4 right);
removing heat from the system causes the equilibrium to replace the removed heat and
the equilibrium therefore shifts right. Added heat shifts the equilibrium in the direction
that absorbs heat; for this reaction, a shift to the left occurs with addition of heat.
This experiment examines the effect of temperature on the system described by
equation 16.17. This system involves an equilibrium between the coordination
spheres, the water versus the Cl– about the cobalt(II) ion; the equilibrium is concentration and temperature dependent. The tetrachlorocobaltate(II) ion, [CoCl4]2–, is more
stable at higher temperatures.
Procedure Overview: A large number of qualitative tests and observations are
performed. The effects that concentration changes and temperature changes have on
a system at equilibrium are observed and interpreted using LeChâtelier’s principle.
The functioning of a buffer system and the effect of a common ion on equilibria are
LeChâtelier’s Principle; Buffers
Ken karp
Figure 16.4 NO2, a red-brown
gas (left), is favored at higher
temperatures; N2O4, a colorless
gas (right), is favored at lower
temperatures. See equation 16.1.
Perform this experiment with a partner. At each circled superscript 1–21 in the procedure, stop and record your observations on the Report Sheet. Discuss your observations with your lab partner and instructor. Account for the changes in appearance of the
solution after each addition in terms of LeChâtelier’s principle.
Ask your instructor which parts of the Experimental Procedure are to be completed. Prepare a hot water bath for Part E.
1. Formation of metal–ammonia ions. Place ~1 mL (<20 drops) of 0.1 M CuSO4 (or 0.1 M NiCl2) in a small, clean test tube. 1 Add drops of conc NH3 (Caution: strong odor, do not inhale) until a color change occurs and the solution is clear (not colorless). 2 2. Shift of equilibrium. Add drops of 1 M HCl until the color again changes. 3 A. Metal-Ammonia Ions 1. Silver carbonate equilibrium. In a 150-mm test tube (Figure 16.5) add ~1⁄2 mL (≤10 drops) of 0.01 M AgNO3 to ~1⁄2 mL of 0.1 M Na2CO3. 4 Add drops of 6 M HNO3 (Caution: 6 M HNO3 reacts with the skin!) to the precipitate until evidence of a chemical change occurs. 5 2. Silver chloride equilibrium. To the clear solution from Part B.1, add ~5 drops of 0.1 M HCl. 6 Add drops of conc NH3 (Caution! avoid breathing vapors and avoid skin contact) until evidence of a chemical change.* 7 Reacidify the solution with 6 M HNO3 (Caution!) and record your observations.8 What happens if excess conc NH3 is again added? Try it. 9 3. Silver iodide equilibrium. After trying it, add drops of 0.1 M KI.10 4. Silver sulfide equilibrium. To the mixture from Part B.3, add drops of 0.1 M Na2S† until evidence of chemical change has occurred. 11 B. Multiple Equilibria with the Silver Ion *At this point, the solution should be “clear and colorless.” † The Na2S solution should be freshly prepared. Figure 16.5 Sequence of added reagents for the study of silver ion equilibria. Experiment 16 211 Disposal: Dispose of the waste silver salt solutions in the Waste Silver Salts container. CLEANUP: Rinse the test tube twice with tap water and discard in the Waste Silver Salts container. Rinse twice with deionized water and discard in the sink. C. A Buffer System 1 2 A Buffer + HCI Buffer + NaOH B H2O + HCI H2O + NaOH The use of a well plate is recommended. Appropriately labeled 75-mm test tubes are equally useful for performing the experiments. 1. Preparation of buffered and unbuffered systems. Transfer 10 drops of 0.10 M CH3COOH to wells A1 and A2 of a 24-well plate, (or appropriately labeled 75-mm test tubes), add 3 drops of universal indicator,† and note the color. 12 Compare the color of the solution with the pH color chart for the universal indicator. 12 Now add 10 drops of 0.10 M NaCH3CO2 to each well. 13 Place 20 drops of deionized water into wells B1 and B2 and add 3 drops of universal indicator. 14 2. Effect of strong acid. Add 5–6 drops of 0.10 M HCl to wells A1 and B1, estimate the pH, and record each pH change.15 3. Effect of strong base. Add 5–6 drops of 0.10 M NaOH to wells A2 and B2, estimate the pH, and record each pH change.16 4. Effect of a buffer system. Explain the observed pH change for a buffered system (as compared with an unbuffered system) when a strong acid or strong base is added to it. 17 D. [Co(H 2 O) 6 ] 2+ , [CoCl 4 ] 2– Equilibrium (Common-Ion Effect) 1. Effect of concentrated HCl. Place ~10 drops of 1.0 M CoCl2 in a 75-mm test tube. 18 Add drops of conc HCl (Caution: Avoid inhalation and skin contact) until a color change occurs.19 Slowly add water to the system and stir. 20 E. [Co(H 2 O) 6 ] 2+ , [CoCl 4 ] 2– Equilibrium (Temperature Effect) 1. What does heat do? Place ~1.0 mL of 1.0 M CoCl2 in a 75-mm test tube into the boiling water bath. Compare the color of the hot solution with that of the original cool solution. 21 Disposal for Parts A, C, D, and E: Dispose of the waste solutions in the Waste Salt Solutions container. CLEANUP: Rinse the test tubes and 24-well plate twice with tap water and discard in the Waste Salt Solutions container. Do two final rinses with deionized water and discard in the sink. The Next Step Buffers are vital to biochemical systems. (1) What is the pH of blood and what are the blood buffers that maintain that pH? (2) Natural waters (rivers, oceans, etc.) are buffered for the existence of plant and animal life (Experiment 20 ). What are those buffers? Experimentally, see how they resist pH changes with the additions of strong acid and/or strong base. (3) Equilibria also account for the existence of hard waters (Experiment 21 ). † pH indicator paper may be substituted for the universal indicator to measure the pH of the solutions. 212 LeChâtelier’s Principle; Buffers Experiment 16 Prelaboratory Assignment LeChâtelier’s Principle; Buffers Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________ 1. a. Describe the dynamic equilibrium that exists between the two water tanks at right. b. Explain how LeChâtelier’s principle applies when the faucet on the right tank is opened. c. Explain how LeChâtelier’s principle applies when water is added to the right tank. 2. a. Experimental Procedure. Cite the reason for each of the five cautions in the experiment. b. Experimental Procedure, Part E. How is “bumping” avoided in the preparation of a hot water bath? 3. The following chemical equilibria are studied in this experiment. To become familiar with their behavior, indicate the direction, left or right, of the equilibrium shift when the accompanying stress is applied to the system. a. NH3(aq) is added to Ag+(aq) + Cl–(aq) b. HNO3(aq) is added to Ag2CO3(s) AgCl(s) _______________ Ag+(aq) + CO3 2– (aq) _______________ c. KI(aq) is added to Ag+(aq) + 2 NH3(aq) d. Na2S(aq) is added to AgI(s) [Ag(NH3)2]+(aq) Ag+(aq) + I–(aq) e. KOH(aq) is added to CH3COOH(aq) + H2O(l) f. HCl(aq) is added to 4 Cl–(aq) + Co(H2O)6 2+ (aq) _______________ _______________ H3O+(aq) + CH3CO2 – (aq) _______________ CoCl4 2– (aq) + 6 H2O(l) _______________ Experiment 16 213 4. Note the dynamic equilibrium in the opening photo. Which solution changes color when the pH of both solutions is increased? Explain. 5. Experimental Procedure, Part C.1. Will the addition of NaC2H3O2 to a CH3COOH solution cause the pH to increase or decrease? Explain. See equation 16.14. 6. A state of dynamic equilibrium, Ag2CO3(s) 2Ag+(aq) + CO3 2–(aq), exists in solution. a. What shift, if any, occurs in the equilibrium if more Ag2CO3(s) is added to the system? b. What shift, if any, occurs in the equilibrium if AgNO3(aq) is added to the system? c. After water is added to the system and equilibrium is reestablished: (i) what change in the number of moles of Ag+(aq) occurs in the system? Explain. (ii) what change in the concentration of Ag+(aq) occurs in the system? Explain. *d. What shift occurs in the equilibrium if HCl(aq) is added to the system? Explain. 214 LeChâtelier’s Principle; Buffers ... Purchase answer to see full attachment

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