KWAME NKRUM AH UN IV ERSITY ERSITY OF SCIENCE SC IENCE AND AND TECHNOLOGY
DEPAR DEPARTM TM ENT OF CHEMI STRY Y EA R T WO (C H EM 2 6 9 ) PRACTICAL PRACTICAL CHEMI CHEMI STRY II I
NAME: OPOKU ERNEST EMAIL:
[email protected]
DATE: 28TH OCTOBER, 2013
A N AL Y SI S O F A M I X T U R E OF C AR BO N AT E A N D BICARBONATE
AIMS AND OBJECTIVES : 1. To introduce titration as a useful technique in chemical analysis. 2. To analyze a sample for total alkalinity and then determine the individual amount of carbonate and bicarbonate.
T H E OR Y AN D I N T R O DU C T I O N Carbonat e a nd Bicarbonate The carbon dioxide that is dissolved by naturally circulating waters appears in chemical analysis principally as bicarbonate and carbonate ions. Carbonate that follows this path represents a linkage between the carbon cycle and the hydrologic cycle. The large supply of atmospheric carbon dioxide is partly intercepted by photosynthesizing vegetation. They convert it to c ellulose starch and related carbohydrates. These products are later reduced via respiration to carbon d ioxide and water with a release of stored energy. The concentration of carbonates in natural waters is a function of dissolved carbon dioxide, temperature, pH, cations and other dissolved salts. At one atmosphere of pressure, pure C02 gas ov er distilled water will produce a solution that would have a pH near 3.6. In the presence of excess calcite, however, the solution would contain some 350 mg/L of dissolved calcium and its pH would be near 6. A considerable part of the released carbon diox ide must return directly to the atmosphere and obviously the efficiency of utilization of the dissolved portion will be far below the theoretical maximum. Water quality records for streams in the United States indicate that the maximum rate for calcium and bicarbonate removal is near 400 tons per square mile per year, but most streams carry far less than half this much. These figures do suggest that under most favorable conditions, limestone may be rather rapidly eroded. It may also be of interest to note that a considerable part of the anionic load of man y streams is a contribution from carbon dioxide of the atmosphere rather than from the rocks of the drainage basin.
Bicarbonate concentration of natural waters generally is held within a mod erate range by the effects of the carbonate equilibria. Most surface streams contain less than 200 mg/L Carbonate and Bicarbona te, but in ground water somewhat higher concentrations are not uncommon. Concentrations over 1,000 mg/L sometimes occur in waters which are low in calcium and magnesium and especially where processes releasing carbon dioxide such as sulfate reduction are occurring in the ground water reservoir. Many of the carbonates are quite insoluble in water, generally more so than the chlorides, nitrates or sulfates. There is a tendency for certain carbonate salts to be removed by precipitation or absorption. In more calcareous environments, the circulation of water rich in carbon dioxide may produce solutions that are highly supersaturated when exposed to the air. Such solutions may deposit large quantities of calcium carbonate as travertine near their points of discharge. In hard waters, particularly in headwaters that are fed b y limestone springs, deposits of calcium carbonate are often layed down. These may form large solid structures that can dam up a stream or produce waterfalls. This precipitated material is travertine. Some of the deposition is probably pu rely chemical and is caused by the loss of equilibrium carbon dioxide nec essary to keep the calcium carbonate and bicarbonate in solution. However, it is nearly always associated with algae and to a lesser extent with mosses that cause deposition of calcium carbonate by photosynthesis. As will be discussed in alkalinity, it is important to note that bicarbonates tend to reach
equilibrium with the carbonates. For this reason, most running waters are “bicarbonate” waters in a limnological sense and show the complicated relationships between pH, C02, H2CO3, H+, C03 – , HCO3-, calcium +2 and magnesium +2.
For this discussion, it is sufficient to stress only a few points: (1) rainwater reaching the water courses as a runoff from bogs, dense forest litter, and similar substrata tends to have a low pH because of the hydrogen ions produced by disassociation of carbonic acid and the loss of cations b y base exchange with the organic matter, (2) water which has percolated through the soil is also rich in carbon dioxide and similarly tends to be rich in hydrogen ions according to this equation: H20 + C02 -> H2CO3 H + HCO3 (3) calcium carbonate which is a common constituent of many rocks is almost insoluble in water, but it dissolves fairly readily as bicarbonate in carbonic acid, and it neutralizes soil water where it occurs, according to this reaction: CACO3 + H2CO3 -> cA(Hco3>2 CA+2 + 2HC03-2. (4) Calcium bicarbonate in solution is a good buffer s ystem and thus resists change in pH, but it remains in solution only in the presence of a certain amount of free carbon dioxide. Any process which removes carbon dioxide, as does photosynthesis, tends to cause precipitation of calcium carbonate from solution, especially where the bicarbonate is abundant.
Therefore, springs in limestone regions are often very rich in calcium bicarbonate where they emerge to the surface. As the water flows along, it looses carbon dioxide to the atmosphere and by photosynthesis. After some distance this loss becomes the loss of equilibrium C02 and the deposition of calcium carbonate occurs by a reversal of the first part of the equation under #3 above. This process is also aided by the ability o f many plants to make direct use of bicarbonate ions as shown in this equation: CA(HCO3)2 -> CAC03 + H20 + C02 (photosynthesized). Deposition of calcium carbonate is therefore a common feature of streams in limestone areas and is a subject which we shall discuss further in the alkalinity section. The carbon dioxide released within the soil by respiration and decay is capable of producing low pH in circulating water if minerals that act as proton acceptors are scarce. Soils of humid, temperate regions may become depleted in calcium carbonate by leaching and the pH of ground water at shallow depths may be rather low. In general, it may be expected that carbonates in themselves are not detrimental to fish life, but their buffering action and effect upon pH ma y contribute to the toxicity of high pH values. Little is known about the influence of carbonate or bicarbonate on higher plants in rivers. Carbon dioxide may be a limiting factor in soft waters where no bicarbonate ions are available because water in equilibrium with normal air containing 4.0 ml/L contains o nly 0.68 ml/L in solution of carbon dioxide. This is much less than is available to land plants. This difference is one of the chief reasons for many of the peculiarities of submerged aquatic plants. To summarize carbonate and bicarbonate, if a spring is from limestone or some other very calcareous rock, the water will be heavily charged with calcium bicarbonate. It will loose carbon dioxide very rapidly and its pH will rise. After some distance, the loss of c arbon dioxide to the atmosphere and by photosynthesis will lead to the depositio n of calcium carbonate. This process will decline steadily as equilibrium is attained. In this well buffered hard water the pH will not rise above about 8.3 even at times of very active photosynthesis, and these changes will occur while the water flows onl y a very short distance from a few hundred meters to a few kilometers, according to the situation. Water from noncalcareous springs will similarly loose carbon dioxide, increase its pH and acquire oxygen in quite a short distance. If the source of the spring is a big, acid swamp, swampy woodland, or rain forest, the water may contain ferrous bicarbonate. As pH rises and oxygen is acquired, ferric hydroxide will be deposited probably over a distance measurable only in tens or hundreds of meters. Unless the water remains acidic, little iron will remain in solution. This part of the stream could be coated with rust colored masses of iron bacteria. If this water is soft, its pH will fluctuate markedly because of photosynthesis and it may exceed 8.3 in the daytime
. The component in a solution containing sodium carbonate and sodium hydrogen carbonate provides examples of how neutralization titration can be used to analyze mixtures. Numerous inorganic species can be determined by titration with strong acids or strong bases. Some of these examples are the mixtures containing hydroxide (OH-), carbonate (CO32-) and bicarbonate (HCO3-) ions. The quantitative and qualitative determination of more than two of these components mentioned above can exist in a substantial amount in any solution, because reaction will remove one of them. Carbonate (CO32-) and bicarbonate (HCO3-) ions requires two titrations. One with a base range such as phenolphthalein and the other with acid range indicator such as The total alkalinity (defined as total base concentration in this case is measured by titrating the mixture with standard HCl to a Bromocresol green end point:
Total alkalinity= [HCO3-] +2[CO32-]. For these samples, this is approximately the amount of
base that reacts with an added acid, hence the term “total alkalinity” and the original amount of carbonate and bicarbonate present in your sample. To determine total alkalinity, you have to titrate your sample with the standard hydrochloric acid using a bromocresol green end point. The reactions are: HCO3-(aq) + H+(aq)→ H2CO3 (aq) CO32-(aq) + 2H+(aq) → H2CO3(aq) In this first titration, there is no way to differentiate between the second and first reactions, so the endpoint of the reaction is for both of the reaction s, together. In the second step, you will determine the original carbonate concentration of your unknown by exploiting a number of methods, including selective precipitation. You will treat a fresh aliquot with an excess of the standard sodium hydroxide from Experiment 1, converting all HCO3- to CO32-: HCO3-(aq) + OH-(aq) → CO32-(aq) + H2O The sample then consists entirely of CO32, which you will precipitate by adding excess barium chloride: Ba2+(aq) + CO32-(aq) → BaCO3(s) The excess NaOH is then immediately titrated with standard hydrochloric acid to determine how much bicarbonate had been originally present in the sample. From the total alkalinity, and bicarbonate the original carbonate concentration can be calculated.
CHEMICALS AND EQUI PMENT 1. Unknown sample 2. Distilled water 3. 0.1M Hydrochloric acid
4. 10%(wt/wt) BaCl2 5. Bromocresol green indicator 6. 0.1M NaOH 7. Phenolphthalein indicator 8. Two 250ml volumetric flasks 9. 25ml Pipette 10. three 250ml conical flasks 11. Funnel 12. Electronic balance, 13. Burette 14. 10ml measuring cylinder
PROCEDURE PROCEDURE 2.01g of the unknown sample was measured into a beaker. The unknown was dissolved in distilled water in the beaker and then transferred into a 250ml volumetric flask and topped up to the mark. An indicator blank was prepared by adding some of the indicator to an amount of BaCl2 and NaOH. The unknown sample solution was then titrated with standard 0.1M HCl. The experiment was repeated for two more values. 25ml of the unknown sample was measured into a 250ml conical flask and then 50.0ml of standard NaOH and 10ml of 10% (wt/wt) BaCl2 were added. Two drops of phenolphthalein indicator was then added. The solution was then titrated with standard 0.1M HCl. The step was repeated for two more values.
OBSERVATION
A colorless solution of the mixture was formed. There was a color change from purple to white solution. This indicated the end point color. There was a color change from light blue to light yellow. There was a color change from light blue to light yellow. A purple solution was formed with some precipitates at the bottom.
There was a color change from purple to white solution with some precipitate at the bottom. There was a color change from purple to white solution with some precipitate at the bottom.
T A BU L AT I O N O F R ES U LT S T I T R A T I ON OF U N K N O W N A G A I N S T H C l
Indicator used: Bromocresol green Number of drops: 2 drops Colour change: Blue to pale yellow
HCO3
H
2H
2 CO3
H 2 CO3
H 2 CO3
Burette reading/mL 1 2 3
Final reading/mL 47.20 27.20 46.90
Initial reading/Ml 20.00 0.10 19.9
Titre value/mL 27.20 27.10 27.00
Average titre = (27.20 +27.10 +27.00) ml = 27.10mL 3 Hence volume of HCl that reacted with the mixture = 27.10mL
T I T R A T I O N O F E X C E SS N a O H A G A I N S T H C l
Indicator used: phenolphthalein Number of drops used: 2 Colour change: pink to white Burette reading/mL 1 2 3
Final reading/mL 29.30 29.70 29.50
Initial reading/mL 0.00 0.00 0.00
Average titre value = (29.30 + 29.70+29.50)ml = 29.50ml 3 volume of HCl that reacted with excess NaOH =29.50ml Equation of reaction between NaOH and the carbonate HCO-3 + OHCO2-3 + H2O ------ 3 the equation of reaction between BaCl2 and the bicarbonate Ba2+ + CO2-3 BaCO3 ------ 4 the equation of reaction between the excess NaOH and HCl is NaOH + HClNaCl + H2O ------ 5
Titre value/mL 29.30 29.70 29.50
CALCULATION AND EVALUATION OF DATA From the equation of reaction between HCl and the mixture total alkalinity = [HCO-3] + 2[CO2-3] 1. 27.20ml of HCl was required to react with the first 25ml of the mixture let x= mol of CO2-3 and let y= mol of HCO-3 in the 25ml aliquot total alkalinity= mol of H+ needed to reach the Bromocresol green end point hence total alkalinity= 2x + y = 0.1 × 27.20 2x + y= 2.72ml 50ml of 0.1M NaOH was added, hence mol of NaOH= 50× 0.1= 5.0mol when 29.30ml of the HCl reacted with the excess NaOH mol of HCl= 0.1 × 29.30= 2.93mol mol of NaOH that reacted with HCO-3= 5.00−2.93= 2.07mol hence 2x= 2.72 – 2.07= 0.65mol x= 0.65/2=0.33mol total alkalinity= 0.65 + 2.07= 0.1088M 25ml 2. 27.10ml of HCl was required to react with the second 25ml of the mixture 2x + y= 2.71mol when 29.70ml of HCl reacted with excess NaOH mol of HCl= 0.1 × 29.70= 2.97mol mol of NaOH that reacted with HCO-3= 5.00− 2.97= 2.03mol hence 2x= 2.71− 2.03= 0.68mol x= 0.68/2=0.34mol total alkalinity= 0.68+ 2.03= 0.1084M 25ml
3. 27.00ml of HCl was required to react with the third 25ml of the mixture, 2x +y=2.70 When 29.50ml 0f HCl reacted with excess NaOH mol ofHCl=0.1x29.50=2.95 mol of NaOH that rected with HCO-.3 =5.00-2.95=2.05 Hence 2x+y=2.70-2.05=0.65mol X=0.325mol Total alkalinity=
.
.
= 0.1080M
Average total alkalinity (x ̅ )=
.
.
.
= 0.1084M
Let x be the values of total alkalinity
̅) 2 (X-X 1.6× 10-7 0.000 1.6× 10-7 ̅ ) 2=3.2× 10-7 Ƹ(X-X
̅) (X-X 4.0× 10-4 0.00 -4.0× 10-7
X 0.1088 0.1084 0.1080
Standard deviation=
∑
standard deviation= √
. ×
[
]
=4.0× 10
The bicarbonate concentrations [HCO2-3] for the three values .
.
.
= 0.0828 = 0.0812 =0.0820
Mean [HCO-3] (y̅) =
.
.
.
= 0.0820
Let y be the values of concentration of bicarbonate (y-y̅) 2 6.4× 10-7 6.4× 10-7 0.00 Ƹ(y-y̅) 2=1.28× 10-6
(y-y̅) 8.0×10-4 -8.0×10-4 0.00
Y 0.0828 0.0812 0.0820
3 the standard deviation of [HCO-3]= ∑
[
]
=√
.
×
= 8.0× 10
The three values for the carbonate concentration[CO2-3] are as follows
.
.
.
=0.0132 =0.0136 =0.0130
Mean [CO2-3]=0.0132 + 0.0136 + 0.0130= 0.0133M 3 X 0.0132 0.0136 0.0130
2 (x-x ̅) 1.0× 10-8 9.0× 10-8 9.0× 10-8 Ƹ(x-x ̅) 2=1.90× 10-7
x-x ̅ -1.0× 10-4 3.0× 10-4 -3.0× 10-4
The standard deviation of [CO2-3=
[
∑
The carbonate concentration [CO2-3] =
( .
]
)
.
= √ .
×
=3.08× 10
=0.0792M
Molar mass of K2CO3= 2(39) + 12 + 3(16)= 138g/mol mass of K2CO3 = (0.0792 × 138× 25)/1000= 0.273g mass of sample= 2.50g % by mass of CO2-3= 0.273 × 100= 10.92% 2.5 The bicarbonate concentrations [HCO2-3] =
.
.
.
molar mass of NaHCO3= 23 + 1 + 12 + 16 + 3(16)= 100g/mol mass of NaHCO3= (0.0100× 100× 25)/1000= 0.025g % by mass of HCO2-3= 0.025/2.50 × 100= 1.00%
Hence % of K2CO3= 10.92(±0.00080) wt% % of NaHCO3= 1.00(±0.00308) wt%
=0.0100M
DISCUSION The titration in the first reaction was to establish a stoichiometric ratio between the HCl, the carbonate and the bicarbonate content of the mixture. However, the second titration was a back titration to be used to calculate the amount of each constituent present in the sample. The barium chloride (BaCl 2) was used to mask the carbonate thus, preventing it from interfering the reaction between the NaOH and the bicarbonate (HCO -3), it was BaCl2that formed the precipitate with CO 2-3 at the bottom of the container. The excess NaOH that reacted with the HCl was subtracted from the initial amount of NaOH to find the amount of NaOH that actually reacted with the bicarbonate. Consequently, showing the amount of bicarbonate present in the sample. From eq.3, the bicarbonate reacted with NaOH to produce the carbonate. This implies that in alkaline solution, the bicarbonate behave as a Bronsted-Lowry acid and donates a proton. Contrarily in eq. 1, it behaved as a Bronsted-Lowry base by accepting a proton from HCl. It can therefore be observed that the bicarbonate is an amphoteric substance. The percentage of the carbonate 10.92(±0.00080) wt% was found to be greater than that of the bicarbonate 1.00(±0.00308) wt% in the mixture because the bicarbonate donates a proton to the water which acts as a weak Bronsted-Lowry base thereby p roducing more of the carbonate.
PRECAUTIONS 1 . Parallax error should be avoided when taking readings from the glass wares. 2 . The solution of mixture should be swirled to enable complete formation of mixture. 3 . As soon as the two drops of phenolphthalein was added the solution was titrated against 0.1M HCl. 4 . Blank solution should be used to give a clear indication about the end point of titration.
CONCLUSION The concentrations of carbonate and bicarbonate were determined. It is known at the end of the experiment that The percentage of the carbonate 10.92(±0.00080) wt% was found to be greater than that of the bicarbonate 1.00(±0.00308) wt% in the mixture because the bicarbonate donates a proton to the water which acts as a weak Bronsted-Lowry base thereby producing more of the carbonate.
REFERENCES 1. Gary D. Christian Analytical chemistry sixth edition ( page 75-85) 2. Statistics for Analytical Chemistry, J.C. Miller and J. N. Miller, 3rd edition, page 16-20 3. Vogel’s hand text book of Quantitative chemical analysis ,fifth edition, G.H. Jeffery, J.Bassett, J.Mendham.pg 55-60 4. Zeebee, Richard. 2006. Marine Carbonate Chemistry. The Encyclopedia of the Earth
