Identification of Cations and Anions in a Solution - Lab Report Example

Identification of Cations and Anions in a Solution Chemistry Lab Report March 4, Qualitative analysis and conductometry, which is a measurement of electrolytic conductivity, are the commonly used methods for identifying cations and anions present in unknown solutions. The two methods were employed to achieve the main objectives of this experiment, which included identification of cations, and anions present in an unknown solution. When employing qualitative analysis in determination of identity of unknown solution, similar tests are performed on a known and unknown sample and their characteristics are matched in order to determine the identity of the unknown. In this experiment, the flame test of the unknown solution gave a pink color indicating presence of potassium ions. The chemical reaction between the unknown solution and barium chloride followed by acidification using dilute nitric acid gave a white precipitate indicating presence of sulfate ions, which resulted from formation of barium sulfate that is an insoluble salt. The pH and conductometry tests gave numerical values of 7.90 and 3.031×10^4 µS/cm, respectively. The unknown solution had a pH similar to that of the known while the conductivity value of the unknown was different from that of the known by a small margin (-0.003×10^4 µS/cm). Synthesis of potassium sulfate from a reaction between potassium hydroxide and dilute sulfuric acid gave a percentage yield of 89.70%. Following the results obtained in this experiment, the identity of the unknown was confirmed to be potassium sulfate. Introduction Physical and chemical characteristics can be used to determine the identity of an unknown compound. Physical properties are those that can be determined without changing the composition of the substance. These include color, taste, odor, melting and boiling point, conductivity and density of a substance. Qualitative analysis is the process of determining the composition of chemical substance in matter by conducting various chemical tests (Davis et al., 2005). Most aqueous samples of various salts can be tested for characteristic reactions by reacting them with dilute nitric acid, dilute hydrochloric acid, aqueous barium nitrate/chloride, silver nitrate, ammonium hydroxide or sodium hydroxide. These reagents are used in determining presence of anions (negatively charged particles) in unknown samples, which must be in aqueous states before the tests can be conducted. A number of signs can be used to deduce whether a chemical reaction has taken place some of which include: Color changes, formation of a precipitate (a solid product which ‘falls out of solution’ because it is insoluble in water), disappearance of a precipitate, evolution of a gas (bubbling), creation of heat or cold (Endothermic or Exothermic reactions). In determining the identity of unknown sample, a known compound or solution is subjected to the same tests as those of the unknown. The identity of the unknown is then determined by matching its characteristic reactions with those of the known samples. Anion analysis can be categorized into four groups namely: the acid-volatile group, the barium precipitate group, the silver precipitate group and the soluble group. The acid-volatile group includes sulfide and carbonate ions, these anions form gases that readily evolve from the solution upon addition of strong acids (equation 1 and 2). CO32- + 2H+ H2CO3 ……………1 H2CO3 CO2 + H2O …………….2 Acidification of sulfide ion produces hydrogen sulfide gas (equation 3) S2- + 2H+ H2S……………………3 The Barium precipitate group, which includes the sulfate and phosphate ions form precipitate upon addition of excess Ba2+ ions. A chemical reaction between Ba2+ ions and SO42- results in the formation of barium sulfate, which is an insoluble salt. The ion half equation is shown in equations 4 and 5. Ba2+ + SO42- BaSO4 ……………..4 Ba2+ + PO43- Ba3(PO4)2 ………….5 Presence of sulfate ions can be differentiated from that of phosphate ions in unknown solution since barium phosphate is soluble in dilute hydrochloric acid, while barium sulfate is not, (the precipitate persists upon addition of dilute hydrochloric acid). The silver precipitate group, which comprises of chloride, bromide, iodide and thiocyanide ions form light-colored precipitate with excess Ag+ ions (equation 6). Appearance of these precipitates varies slightly for each anion therefore; one is able to distinguish between them. Ag+ + X- AgX ………………..6, (X= Br-,Cl-, I-, SCN-). The confirmatory test of thiocyanide ions is the blood-red complex it forms with Fe3+ ions (equation 7) Fe3+ + SCN- [Fe(SCN)]2+ ……7 The soluble group consists of nitrate and acetate anions. Nitrate ion is normally identified by the presence of a brown ring which is caused by the formation of Fe(NO)2+ in the presence of NO3-and excess Fe2+ (equations 8 and 9). 3Fe2+(aq) + NO3-(aq) + 4H+(aq) 3Fe3+(aq) + NO(aq) + 2H2O ………..8 NO(aq)+ Fe2+(aq) Fe(NO)2+(aq) (brown)……………………………….9 The H+ ion in equation 8 is sourced from concentrated sulfuric acid. Sulfuric acid forms a lower layer when added to an aqueous solution due to its higher density. The solutions are layered rather than mixed because the heat of dilution of sulfuric acid is enough to destroy the brown Fe complex formed. The brown ring forms at the interface between the two layers (Davis et al., 2005). In cation analysis, the use of flame test is one way in which identity or possible identity of a metal ion found in an ionic compound can be determined. In this test, a nichrome wire is dipped into a solution containing the unknown cation and immediately the wire is held the in the hot burner flame, which gives a characteristic color that is visible to the naked eye. Each metal ion has a characteristic distinct flame color. Potassium ion gives a purple or pink (lilac) flame sodium ion gives a yellow flame color, lithium gives a crimson flame color, calcium gives a brick-red flame color while copper gives a green flame color with blue streaks when held in a hot burner flame. Since each ion is unique and reacts in its own way, conducting the right tests while following the rules of logic helps to determine the identities of the ions present in an unknown solution. In this experiment, qualitative analysis and conductometric analysis were employed in determining the identity of a solution containing unknown cations and anions. Similar tests were conducted on an unknown substance and a known substance in order to determine the unknown substance’s identity by matching its characteristic reactions with those of the known substance. The experiment emphasized more on the importance of accurate observations made, proper recording, and the use of deductive logic to draw conclusions from those observations. Procedure 1.1 Qualitative analysis A solution of 50 ml of unknown solute was prepared by dissolving 0.5 g of unknown solute into 50 ml of distilled water in a clean beaker. A Nichrome wire with a loop was the dipped in the unknown solution and immediately placed in the hot Bunsen burner flame in order to note the color of the flame. In anion analysis, presence of carbonate ion in the solution was tested by adding 1.0 ml of 6.0 M hydrochloric acid measured using a graduated cylinder onto the unknown solution contained in a clean beaker. After confirming absence of carbonate ions in the solution, a solution of 1.0 ml of 0.1 M barium chloride was then added to unknown solution in order, determine the presence of sulfate ions. 1.2 pH determination Calibration of the pH meter was done by placing the electrode into the buffer solution of pH 7 and setting the pH meter to the pH 7 at the measured temperature. The same steps were repeated with a buffer of pH 10 before the pH of the unknown solution was determined in order to get an accurate and repeatable measurement. A conductivity probe and Logger pro were used to determine the conductivity of the unknown solution and the known potassium sulfate and a comparison was made between the two numerical values in order to ascertain the identity of the unknown. 1.3 Preparation of potassium sulfate salt The mass of a dry clean beaker was determined using an analytical weighing balance. In the weighed clean beaker, 11.48 ml of 1.0 M potassium hydroxide measured using a graduated cylinder was poured followed by 5.378 ml of 1.0 M sulfuric acid. The mixture was then stirred using a glass rod in order to increase the kinetic energy of the reacting molecules and hence speed up the rate of the chemical reaction for formation potassium sulfate salt. The mass of the synthesized potassium sulfate salt was obtained by weighing the mass of beaker and its content and subtracting the mass of the empty beaker from that it. The percentage yield of the prepared potassium sulfate salt was determined using the actual mass of potassium sulfate formed and the theoretical mass of potassium sulfate generated from the balanced chemical equation. Results 1.1 Qualitative analysis Test observation inferences Flame test Pink flame color K+ ions present Addition of dilute hydrochloric acid (HCl) No precipitate formed CO32- ions absent Addition of Barium chloride (BaCl2) White precipitate formed SO42- ions present 1.2 pH and conductivity test Known (Potassium sulfate) Unknown Molar mass (g/mol) 174.27 pH 7.90 7.90 Mean Conductivity (×10^4 µS/cm) 3.028 3.031 1.3 Preparation of potassium sulfate Using the balanced chemical equation between potassium hydroxide and dilute sulfuric acid 2KOH + H2SO4 K2SO4 + 2H2O Mole ratio of H2SO4 : K2SO4 = 1:1 Moles of H2SO4 that reacted with KOH to form K2SO4 = volume of sulfuric acid (L) × molarity of sulfuric acid: 0.005378 L × 1.0 M = 0.005378 = moles of K2SO4 formed Mass of K2SO4 formed = moles × molar mass = 0.005378 mol × 174.27 g/mol = 0.99996g ≈ 1.0 g Therefore, 1.0 g is the theoretical yield of potassium sulfate that can be formed based on mole ratio of 1:1 (sulfuric acid : potassium sulfate) from a well-balanced chemical equation while 0.0897 g is the actual yield of potassium sulfate that was formed from the reaction between potassium hydroxide and dilute sulfuric acid. Percentage yield of potassium sulfate salt = (Actual yield × 100%)/ Theoretical yield = (0.897 g × 100%)/1.0 g = 89.70 % Discussion From the results obtained in section 1.1, the pink flame color from flame test confirms presence of K+ ions in the unknown compound. Purple-pink commonly known as lilac flame color is used as a confirmatory test for K+ ions (Practical chemistry, n.d.). In anion analysis, lack of precipitate formation on addition of hydrochloric acid to the unknown solution confirms the absence of carbonate ions in the solution. Formation of white precipitate upon addition of barium chloride confirms presence of sulfate ions in the unknown solution. A chemical reaction between barium chloride and a solution containing sulfate ions results in the formation of barium sulfate, which is an insoluble white salt hence the white precipitate (Slowinski, Wolsey, & Rossi, 2011). Based on the observations made from the qualitative analysis, the unknown solution contains K+ ions and SO42- ions. In the conductometric analysis, the conductivity of the product formed when potassium hydroxide reacts with dilute sulfuric acid is measured due to the change in concentration of the ions present in a solution and their inherent ability to conduct electricity (equivalent conductance). Any reaction that changes the concentration of ions in an ionic solution will show changes in conductivity of that solution. Different salts have different electrical conductivity at standard conditions since each ion has its own specific ability to conduct current, which makes it easy for identification of ions in a solution. The numerical value of electrical conductivity of the unknown solution differs slightly from that of the known (potassium sulfate) by -0.003×10^4 µS/cm. This difference might have been resulted from change in temperature of the solution and storage time (time between preparation of the solution and the testing time), which affect electrical conductivity of a substance. The electrical conductivity measurement should therefore be made either in situ or immediately after preparation of the solution. A chemical reaction between a base and an acid produces a salt and water as the only products, a process called neutralization. Therefore, a reaction between potassium hydroxide and dilute sulfuric acid results in the formation of potassium sulfate, a salt, and water. Using the mole ratio of sulfuric acid to that of potassium sulfate, one is able to calculate the moles of potassium sulfate and hence the theoretical yield of potassium sulfate salt prepared. The percentage yield of potassium sulfate prepared in this experiment was 89.70%, which is less than 100%. The difference might have been due to factors such as impurities present in the reagents used. Conclusion The flame test conducted on the unknown solution in this experiment gave a pink flame color confirming presence of K+ ions in the unknown while the barium precipitate test, which gave white precipitate, indicated presence of sulfate ions in the unknown solution. The unknown solution had a pH of 7.90 and an electrical conductivity of 3.031×10^4 µS/cm while the known solution had a pH of 7.90 and an electrical conductivity of 3.028×10^4 µS/cm. The difference in electrical conductivity between the known and the unknown was 0.003×10^4 µS/cm. The identity of the unknown was confirmed to be potassium sulfate based on the results of the flame test, pH test and the conductivity tests that were conducted. The unknown solution therefore contained potassium ions and sulfate ions. The percentage yield of potassium sulfate synthesized was 89.70 %. The discrepancies observed in these results might have been because of various environmental conditions such as changes in the room temperature and presence of impurities in the final product. References Davis, R. E., Frey, R., Sarquis, M., Sarquis, J. L. (2005). Modern Chemistry. Houghton Mifflin School. Practical chemistry. (n.d.). Retrieved from http://www.practicalchemistry.org/flame-colours-a-demonstration,102,Ex.html. Slowinski, E., Wolsey, W., & Rossi, R. (2011). Chemical principles in the laboratory. Belmont, CA: Brookes & Cole. Read More