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Chemistry practical ppt

  3. 3. BACKGROUND INFORMATION Important skills required in chemistry practical Practical is meant to assist the learners to learn and develop the following skills; • Manipulative skills • Observation skills • Reading skills • Recording skills • Computing skills • Interpreting skills
  4. 4. • Manipulative skills- involve the correct and skillful handling of the apparatus by the experimenter. • Observation skills- involve the use of sense: Hearing, sight, touch and smell to detect changes/reactions. • Reading skills- refer to the ability to comprehend written or printed information. • Recording skills- entails writing down, for example, measurements or statements of facts or other details for reference. • Computation skills- involve the use and application of mathematical knowledge on the data collected. • Interpretation skills- involve studying the data collected and drawing conclusions based on the established chemistry principles.
  5. 5. Areas to consider when preparing for KCSE Chemistry • Chemical symbols must be written correctly i.e. • When the chemical symbol is just one letter, it must be in capital. • If two letters are used the first must be a capital and the second one be a small letter. • They must be separate (not joined). • The small letter must not be bigger than the capital ones. • Must not be shaded.
  6. 6. •In the chemical formulae, any letter that represents one atom of an element must be capital e.g. Na2CO3 and not Na2Co3. •A chemical equation MUST contain correct symbols and formulae, balanced and state symbols to show the state of mater of each of the substances involved.
  7. 7. Writing observations must be based on the following: • When there is colour change, both original and final colours MUST be stated. • Colours and smell of gases should be stated. • Colours of precipitates should be stated. • Colour changes for both blue and red litmus papers MUST be stated when used. • Colours of solutions must be stated.
  8. 8. A procedure must be correct to the letter e.g. • Soluble salts must be prepared by adding EXCESS metal to an acid, metal oxide to an acid or metal carbonate to an acid. • In separating a soluble solid from an insoluble one, start by adding a convenient solvent to the mixture. DO NOT say dissolve the mixture. • Insoluble salts are prepared by precipitation (mixing two solutions). Indicate what happens at each step.
  9. 9. When naming or identifying substances, do it as follows; • When the questions leads identify, either the name or formula is allowed. • But when the question leads name, then just give the exact name. • Names of ions are not allowed only correct formulae are accepted. • If you give both the name and the formula, both MUST be correct.
  10. 10. When plotting graphs, ensure they have the following: • Labeled axes. • Appropriate scale. • Plotted points must be visible. • All points in the table must be plotted correctly. • A convenient scale must make the graph occupy at least ¾ page. • Straight lines must be drawn with a ruler. • Smooth curves must be drawn free hand.
  11. 11. Drawing diagrams • Must be correctly drawn • Apparatus should be proportional • Labeling • Must be workable i.e no lickage of the gas
  12. 12. Conditions under which the reactions were carried out must be stated e.g. • Catalyst • Specific temperature or range of temperature • Specific pressure • Specific concentrations. NB: Abbreviations are not accepted e.g. ppt instead of precipitate
  13. 13. Filling the Burette: • First rinse with about 5cm3 of the solution with which it is to be filled. • Discard the portion of solution used to rinse it. • Fill the burette with the solution to the mark. The level of the bottom of the meniscus is then adjusted to be at the same level with the zero mark. This is taken as the initial reading. • Any convenient mark may be chosen as the initial reading for any other subsequent filling
  14. 14. • Writing ions in words and wrong oxidation states is not acceptable. Ions should be written in their symbols and correct oxidation states e.g. Zn2+ and not Zn+2. • When mixing two solutions it is wrong to write: • No reaction • No change • No chemical change/reaction • Clear solution formed.
  15. 15. • Correct observation should be: • A white precipitate formed. • A colured (mention the colour) precipitate is formed. • No white (or no coloured) precipitate is formed
  16. 16. • There is no white nor whitish solution. A solution is colourless if no colour is seen in the solution. • Do not add-ish to colours(s) of precipitate of solutions e.g. copper solution is blue not bluish.
  17. 17. Using volumetric flask when: a) Diluting solutions • Measure the required volume of solution to be diluted and transfer it to the flask. • Add more distilled water until the bottom of the meniscus is at the level of the mark on the neck. b) Preparing solutions from solid substances. • Dissolve the solid in a little water in a beaker. • Transfer the solution into the volumetric flask. • Add more distilled water until the bottom of the meniscus is at the level of the mark on the neck of the volumetric flask.
  18. 18. VOLUMETRIC ANALYSIS • This is a specialty in chemistry which deals with measurements of volumes of solutions of reactants, reacting them as you record the reacted volumes and using the information to carry out calculations based on the data collected.
  19. 19. Terms associated with volumetric analysis • Standard solution:- is a solution of known concentration which is used in titrimetry. • Analyte:- is the test solution whose concentration is to be determined. • Equivalent volumes: are the volumes of both the standard and analyte solutions which react completely without any in excess. The end of the reaction is known as the end point or equivalent point.
  20. 20. Commonly used end point indicators • For neutralization reactions, phenolphthalein and methyl orange are used. • Redox titrations do not require an indicator since they change colour when reduced during the reaction e.g. • K2MnO4- (Potassium manganate (Vii) changes colour from purple to colourless). • K2Cr2O4- (Potassium dichromate (Vi) changes colour from orange to green).
  21. 21. Indicator Colour in; Acid Base Neutral solution Methyl-orange Screened methyl orange Red/Pink Purple Yellow Orange Orange Orange Phenolphthalein Colourless Pink Colourless Litmus solution Red Blue Purple NB: Use 2 or 3 drops of indicator to get precise colour change/end point.
  22. 22. THE MOLE • The mole is the chemist’s unit of measurement of the amount of substances. • It is a measuring unit just like a dozen, a pair, or a gross e.g. • A pair of shoes consists of 2 shoes. • A dozen shoes consist of 12 shoes. • A mole of shoes consist of 6.023 x 1023 shoes.
  23. 23. • In chemistry a mole of a substance is the amount of that substance containing 6.023 x 1023 particles. • In this case, particles can be atoms, ions, molecules or electrons. • The mole of a substance may be represented by the chemical symbol or the chemical formula of the substance referred to e.g.
  24. 24. • Cl- represents a mole of chlorine atoms. • Cl2 – amole of chlorine molecules. • O- a mole of oxygen atoms • O2- a mole of oxygen molecules. • Cu- amole of copper atoms. • Cu+2- a mole of copper ions. • e- a mole of electrons.
  25. 25. MOLAR MASS • This refers to the mass of one mole of the substances. The unit of measurement for the mass is g/mol e.g
  26. 26. Amole of: Symbol or formula Molar mass (g/mol) Oxygen atoms O 16 Oxygen molecules O2 32 Sodium chloride NaCl 58.5 Copper (ii) oxide CuO 79.5 Copper atoms Cu 63.5 Hydrogen Chloride molecule HCl 36.5
  27. 27. Concentration of a solution • This is the amount of a solute dissolved to make a given volume of solution. • The concentration of a solution may be expressed in terms of: • mass per litre of solution (g/l) or (g/dm3) • or • moles per litre of solution (mol/l) or (mol/dm3)
  28. 28. The number of moles in one litre solution is called molarity (M) A solution containing one mole of a substance per litre of solution is referred to as a molar solution. A molar solution, therefore, contains molar mass of the solute per litre of the solution.
  29. 29. Relationship between Molarity, Number of moles, Volume of solution, RFM and Mass per litre of solution • Molarity = Number of moles Volume in litres • Number of moles = Molarity x Volume inlitres • Volume in litres = Number of moles Molarity Mass of one mole = RFM of the substance in grams
  30. 30. When diluting solutions use the formula • C1 V1 = C2 V2 Or • M1V1=M2V2 Where; • C1 is the initial concentration • C2 is the final concentration • V1 is the initial volume • V2 is the final volume • M1 is the initial molarity • M2 is the final molarity
  31. 31. TITRATION • Titration is a process in which a standard solution is carefully reacted with an analyte with the sole aim of determining the concentration. It involves measurement of volumes and calculations. • Tests on learner’s ability to manipulate the burette, pipette and conical flask. • The concentration determined can then be used to find out either; • Relative atomic mass • Percentage composition in the compound used and • Reacted mass • Experiment is repeated 3 times until atleast two values with a deviation of ±0.2 cm3 of each other are obtained (consistent values).
  32. 32. Types of titration • There are several types of titrations. But, only four are commonly carried out at secondary school level. These are: • Direct titrations • Back titrations • Double indicator titrations • Redox titrations.
  33. 33. Direct Titrations These types involve the additions of a standard solution to a fixed volume of the analyte(solution whose concentration is not known) with the aim of determining either; • Its concentration; • The RAM of one of the constituent elements of the sample analyte; or • The percentage composition of the constituent elements.
  34. 34. Back Titration • A back titration involves reacting an analyte with excess of a reactant of known concentration. The excess reactant is then titrated against a standard solution. • A back titration is preferred where a reaction is too slow such that a sharp end point cannot be attained easily or where a good end point indicator cannot be identified.
  35. 35. Double indicator Titration • Double indicator titration is carried out with the aim of determining the composition of a mixture of different compounds. Example • You are provided with a solution X that contains 12.72 g dm-3 of Na2CO3 mixed with an unknown mass of NaHCO3 . 25cm3 of the solution mixture is first titrated with 0.1 M HCl using phenolphthalein indicator. The analyte is further titrated with the acid using methyl orange indicator.
  36. 36. FOR A SUCCESSFUL VOLUMETRIC ANALYSIS CONSIDER DOING THE FOLLOWING: • Complete the table and ensure that your values are consistent i.e. within ±0.2 of each other. • A candidate may be penalized the mark(s) for complete table if: • Arithmetic’s are wrong • Table is inverted i.e. values being misplaced e.g. initial burette reading in the column of final burette reading and vice versa. • Values are above 50cm3 unless explained. • Values are unrealistic e.g. above 100cm3 or below 0.5 cm3
  37. 37. • The candidate should use 1 0r 2 decimal places when filling the table. Uniformity of the decimal is a MUST. • If you use 2 decimal places, the 2nd one must be a zero or a five e.g. 20.10 or 20.15 and can NEVER BE 20.13 or 20.16 etc. • Values to be averages must be within ±0.2 of each other.
  38. 38. • All working must be shown starting with the calculations of the average volume (the principle of averaging) all the way to moles and molarity. • Transfer the values infact i.e. values as they are from the preceding questions. • Use correct units. Wrong units call for penalty of ½ mk. Better leave it if not sure • It is important to know the colour of indicators in acids and basic solutions.
  39. 39. PHYSICAL CHEMISTRY This area deals with: • Thermo-chemistry/energy changes • Rates of reactions • Solubility • Electro-chemistry
  40. 40. Thermo-chemistry/energy changes • Deals with energy in a system. This is because as chemicals or physical changes take place, a system may give out heat energy to the surrounding or absorb from it. • Any experimenter has to be very keen in collecting accurate data and making correct observations. • Any one given chemical has a certain amount of energy associated with it.
  41. 41. The amount of energy associated with a system is referred to as the enthalpy of the system (symbol H) The amount of energy gained or lost in known as enthalpy change (symbol ΔH) A reaction in which energy is gained is referred to as an endothermic reaction. A reaction in which energy is evolved /gained is said to be an exothermic reaction. NB: THESE TYPES OF CHANGES ARE NOTED SINCE THEY INVOLVE TEMPERATURE RISE OR FALL ON THE THERMOMETOR.
  42. 42. Energy level diagrams An energy LEVEL diagram illustrates the; change in energy, type of reaction, reactants and products in a tabular form. LEVEL A LEVEL B H1 PRODUCTS REACTANT PRODUCTS REACTANTS H2 H2 H1
  43. 43. Where; • H1 is the initial enthalpy • H2 is the final enthalpy • ΔH is the enthalpy change. • In calculations, use the formula; ΔH= H2 – H1
  44. 44. Which energy level diagram represents; •An endothermic reaction? •An exothermic reactions? Answer: •Level A •Level B
  45. 45. Terms associated with enthalpy changes • Heat of Reaction (ΔH reaction): this refers to the amount of heat energy absorbed or evolved when a reactants change to products. • Standard heat of Formation (ΔHf): is the amount of heat energy absorbed or liberated when one mole (1M) of a substance is formed from its constituent elements under standard conditions of temperature and pressure. • Heat of Combustion (ΔHc): is the quantity of heat evolved when a given mass of a chemical substance is burnt in excess oxygen.
  46. 46. •Molar Heat of combustion (ΔHc): is the amount of heat change when one mole of a pure substance is burnt in excess oxygen under standard conditions of temp and pressure. •Latent Heat: this is the amount of heat energy required to cause the physical change of a given mass of a liquid to vapour at constant temperature i.e. heat needed for; melting, freezing, evaporation, or condensation. •Heat of neutralization (Hneut): is the amount of heat energy evolved when one mole of a solution of an acid neutralizes one mole of alkali solution to form an infinitely dilute solution.
  47. 47. •An infinite dilute solution is one which is so much dilute such that there is no further enthalpy change that occurs when more water is added. •Lattice Energy (ΔHlatt): is the amount of energy liberated when isolated gaseous ions combine to form an ionic solid. •Hydration Energy (ΔHhyd): is the amount of heat energy liberated when isolated gaseous ions interact with and are surrounded by water molecules.
  48. 48. Determination of Enthalpy changes. Most reactions in thermo-chemistry occur in aqueous solutions. The following assumptions are always made when working out enthalpy changes. •The density of the solution is always equal to that of water. Which is 1 x 103 kgcm-3 •The specific heat capacity is always the same as that of water. Which is 4.2 KJKg-1K-1 •The solution formed is infinite dilute •No heat is lost to the surrounding or gained from the surrounding during the experiment. •Dissolving a small amount of solid does not significantly alter the final volume of the solution.
  49. 49. the sources of error and account for the differences between experimental and theoretical values. • EXAMPLES • Suppose excess iron is added to 25cm3 of a 0.2M CuSO4 solution at 25⁰c. if the highest temperature reached is 35. Determine the molar enthalpy of the reaction. • It should be assumed that all the copper ions in the 25cm3 of 0.2M CuSO4 are displaced by iron. • Equation for the reaction;
  50. 50. Equation for the reaction; Cu2+(aq) + Fe(s) Cu (s) + Fe2+(aq) Mole of copper displaced = Molarity x volume in litres = 0.2 mol x 25cm3 1000cm3 = 0.005 moles Mole ratio of Fe: Cu = 1:1 Therefore no. of moles of iron oxidized to Iron (II) ions = no. of moles of copper (II) ions reduced to copper = 0.005 moles Amount of heat (Q) = Mass in Kg (M) x specific heat capacity (c) x change in temp (ΔT) Q = MC ΔT
  51. 51. = 25g x 4.2KJ x (35.2 -25) K 1000g x kgk = 0.025 kg x 4.2kj x 10.2 K Kg K = 107.1 KJ This means that 0.005 moles evolved 107.1 Kj. Molar heat of solution is the amount of heat evolved if one was used Which is 107.1 KJ x 1 mole 0.005 mole = 212 Kj/mol
  52. 52. Molar heat of solution is the amount of heat evolved if one mole was used Which is 107.1 KJ x 1 mole 0.005 mole = 212 Kj/mol
  53. 53. Molar Heat of Neutralization of sodium Hydroxide. This can be investigated by determining the change in temperature that occur when different volumes of 2M solution of unknown acid. Record the result in the table. Sample results Experiment I II III IV V VI VII Volume of 2m NaOH (aq) 10 15 20 25 30 35 Volume of 2M Acid x Initial temperature(⁰C) 35 30 25 20 15 10 Final temperature (⁰C) 24.0 24.0 24 .0 24.0 24.0 24.0 Temperature change (⁰C) 29.0 33.0 34 .5 37.5 36.5 33.0 5.0 9.0 10 .5 13.5 12.0 9.0
  54. 54. Questions •Plot a graph of the rise in temperature against the volume of sodium hydroxide solution added. •From the graph determine the greatest temperature change •What is the significance of the greatest temperature change? •Assuming that the density of the resulting solution in experiments is 4.2jkg-1/k-1 calculate the molar heat of neutralization of 2m NaoH (aq).
  55. 55. Experiment to investigate molar Enthalpy of Displacement. •You are provided with: •Solution Y containing 0.2 moles of Copper (II) Sulphate per litre of solution •Solid Z •You are required to determine the heat evolved when one mole of solid Y reacts with solid Z. Procedure: •Measure 40cm3 of solution Y and place it into an insolated 50cm3 plastic beaker. •Stir the solution with the help of the thermometer and record its temperature after every ½ minute for 1 ½ minutes. •After exactly 2 minutes, add all the solid Z provided and continue stirring the mixture while recording the temperature of solution and complete the table 1 below.
  56. 56. Time mins ½ 1 1 ½ 2 2 ½ 3 3 ½ 4 4 ½ 5 5 ½ 6 6 ½ Temp •On the graph paper provided plot a graph of temperature against time. •From your graph, determine the maximum temperature change. •Given that the density of the solution is 1g/cm3, determine the quantity of heat evolved when the 40cm3 of solution Y is reacted completely with solid Z. •Given that solid Z is Zinc powder, write an ionic equation of the reaction which occurs. •Determine the moles of copper ions used up in the reaction.
  57. 57. •Determine the amount of heat that would be evolved if one mole of Copper (II) ions were used up. •State with reasons whether the value obtained in this reaction is higher or lower than the one obtained from data books.
  58. 58. SOLUBILITY AND SOLUBILITY CURVES • Solubility is the maximum amount of a solute which can dissolve in a given mass of a solvent at a given temperature. • A water soluble substance dissolve in water until no more can dissolve at a given temperature. When this happens, the solution is said to be saturated, • The temperature at which a given mass of solvent cannot dissolve any more solute is called saturation temperature.
  59. 59. When such a solution is cooled below its saturation temperature, the solution forms crystals of the solute. Solubility of most solids increases with increase in temperature; however, solubility of a few solids decreases with temperature rise. Solubility is expressed in or mole (mol.) per 100g of the solvent. Solubility of a substance can be determined accurately by experiments in different ways as illustrated below:
  60. 60. Rates of reactions This is also referred to as chemical kinetics. • It involves the investigation on the effects of changing one of the conditions/parameters on the time taken for reactants to be consumed or the products to be formed. Factors affecting the rates of reactions include: • Concentration of reactants • Surface area (partial size) • Pressure • Temperature • Presence or absence of a catalyst
  61. 61. The dependency of a given factor is investigated if the rest are kept constant. Determination of rates of reactions involves measurements of; time, temperature, volume and the mass of reactants or products. For good results the learner must be able to use calibrated instruments accurately. Plotting of graphs using data collected is a common feature in the study of rates of reaction. The following information and skills on graph work are a requirement for the learner; •Giving a graph a concise, accurate and specific title •Choosing and using a suitable scale •Labeling the axes accurately •Plotting the collected data correctly •Joining all plotted points for a curve or the best line of fit for a linear graph
  62. 62. Qualitative Analysis • Qualitative analysis refers to the process of carrying out chemical tests on substances with the sole aim of identifying them. • Involves identification of ions without taking any precise measurements. • It involves close and careful observations that are made on the chemical changes that occur during chemical reactions. • Reasonable conclusions/inferences are made based on the observations e.g. when a substance is dissolved in water, heated, acid is added, NaOH (aq) or NH3 (aq).
  63. 63. Qualitative analysis is divided into two parts: • Part 1: Qualitative inorganic analysis • Part 2: Qualitative organic analysis • In both areas, to obtain good results, a student must be able to: • Accurately identify the various test reagents. • Identify what these reagents tests. • Predict the expected results.
  64. 64. The following are important points to note • Formation of precipitate generally leads to the identification of cations. • Evolution of a gas leads to the identification of anions. • Testing for the presence of cations involves the use of dilute NaOH or ammonia solutions as the preliminary test reagents. • Identification of anions involves the use of dilute HCl, (vi) acid or nitric (v) acid as the preliminary test reagents. • Flame tests involves the use of Nichrome wire or clean metallic spatula or glass rods dipped in a solution containing the metal ions and hating it in a non-luminous flame.
  65. 65. • Colour of residues or oxides of certain metals is also necessary in identifying ions e.g. • Red resides when hot, yellow on cooling indicates presence of Pb+2 as the residue is PbO. • Black reside when hot or cold likely suggests FeO, CuO hence Cu2+, Fe2+ ions might be present but further test needed to eliminate one. • Yellow reside when hot and white on cooling indicates that the residue is ZnO, therefore Zn2+ ions present.
  66. 66. Inorganic Analysis • The basic test for gaseous products include: • Colour • Smell • Effect on limewater • Effect on a glowing or burning splint. • NB: Students often ignore basic tests for gases even when one has reasons to believe a gas is being evolved.
  67. 67. • Identification of a gas being evolved may be one of the clues required to identify the substance being tested. Background information on salts. • Solubility of salts • Effects of heat on salts
  68. 68. Solubility of salts • All nitrates are soluble in water • All chlorides are soluble in water except AgCl and PbCl2.Lead chloride dissolves on warming and crystallizes on cooling. • All carbonates are insoluble in water except Na2CO3, K2CO3 and (NH4)2CO3. • All sulphates are soluble in water except PbSO4, BaSO4 and CaSO4. Calcium sulphate is slightly soluble. • All sulphites are insoluble in water except ammonium and alkali metals sulphites.
  69. 69. • : Barium sulphite dissolves in dilute nitric (V) acid and HCl while Barium sulphate does not. This acts as a test for distinguishing the two salts. THE TABLES BELOW SHOWS THE PHYSICAL NATURE AND EFFECT OF HEAT ON SALTS
  70. 70. Colour of salt Cation present White solids & forms colourless solutions K+, Ca2+,Na+, Mg2+, Al3+, Zn2+, Ba2+,Ag+, Pb2+,NH+ 4 Pale green Fe2+ Green Ni2+ and Cr3+ Yellow/brown Fe3+ Blue/ pale green depending on the compound Cu2+
  71. 71. Nitrate of: Effect of gentle heating Effect of strong heating K and Na Melts to a colourless liquid Decomposes to yield oxygen and nitrites Ca, Mg, Al, Zn, Fe, Pb and Cu Melt to a liquid Decomposes to produce oxygen, nitrogen (Iv) oxide and metal oxides. NH4 + Melts to a colourless liquid Decomposes to nitrogen (i) oxide and water
  72. 72. Carbonate of: Effect of gentle heating: Effect of strong heating: K and Na No effect Do not decompose Ca,Mg,Zn, Pb, Fe and Cu No effect Decomposes to produce carbon (iv) oxide gas and metal oxide. NaHCO3 No effect Decomposes to form carbon (iv) oxide, water and metal carbonate (NH4)2CO3 Slow decomposition Decomposes to form ammonia, carbon (iv) oxide and water
  73. 73. Effect of heat on chlorides Heat has no effect on chlorides. Effect of heat on Sulphates :- Iron (II) Sulphate decomposes on heating liberating mixture of sulphur (IV) and sulphur (VI) oxides and the oxide of the metals. Ammonium sulphate decomposes to ammonia gas and ammonium hydrogen sulphate.
  74. 74. Gas Colour and smell Test Test result Ammonia (NH3) - Colourless - Has a pungent smell Hold moist litmus paper ( or universal indicator paper) in the gas. Indicator turns blue Carbon (IV) Oxide - Colourless - Odourless Bubble the gas through lime water (calcium hydroxide solution) White precipitate of calcium carbonate is formed. Chlorine (Cl2) - Pale green - Has a choking smell Hold moist litmus paper (or universal indicator paper) in the gas. Indicator paper is bleached. (Wet blue litmus will turn red first before getting bleached.) Properties and Tests for Common Gases
  75. 75. Hydrogen (H2) - Colourless - Odourless Hold a burning wooden splint in the gas. Burns with a pop sound. Oxygen (O2) - Colourless - Has a choking smell Hold a glowing wooden splint in gas Relights the splint
  76. 76. • NB: • Do not smell gases directly instead lift it towards the nostrils • Litmus paper used must be moist and should not touch the walls or the test tube. • Both blue and red litmus paper must be used.
  77. 77. Identification of cations • Cations are positively charged ions. They are identified through tests in which precipitates and complex ions are formed when sodium hydroxide or aqueous ammonia is added to their salt solutions. • The method is based on the principle that, insoluble metallic hydroxide are formed when two to three drops of aqueous ammonia or sodium hydroxide are added to the test solution.
  78. 78. • NB: The colour of some of the precipitate and solubility of some of the hydroxides in excess alkali, help to identify the ions present in the substance.
  79. 79. Cati on Reaction with NaOH (aq) Reaction with aqueous ammonia (NH3) Ba2+ - White precipitate formed - Precipitate insoluble in excess alkali - White precipitate formed - Precipitate insoluble in excess aqueous ammonia Ca2+ - White precipitate formed - Precipitate insoluble in excess alkali - White precipitate formed - Precipitate insoluble in excess aqueous ammonia Mg2+ - White precipitate formed - Precipitate insoluble in excess alkali - White precipitate formed - Precipitate insoluble in excess aqueous ammonia
  80. 80. Al3+ - White precipitate formed - Precipitate soluble in excess alkali - White precipitate formed - Precipitate insoluble in excess aqueous ammonia NH+ 4 - No white precipitate - When heated a gas which turns moist red litmus paper blue is evolved - No white precipitate - When heated a gas which turns moist red litmus paper blue is evolved. Zn2+ - White precipitate formed - Precipitate soluble in excess alkali - White precipitate formed - Precipitate soluble in excess aqueous ammonia
  81. 81. Fe2+ - Green precipitate formed - Precipitate insoluble in excess alkali - Green precipitate formed - Precipitate insoluble in excess aqueous ammonia Fe3+ - Brown precipitate formed - Precipitate soluble in excess alkali - Brown precipitate formed - Precipitate insoluble in excess aqueous ammonia Pb2+ - White precipitate formed - Precipitate soluble in excess alkali - White precipitate formed - Precipitate insoluble in excess aqueous ammonia
  82. 82. Cu2+ - Blue precipitate formed - Precipitate insoluble in excess alkali - Blue precipitate formed - Precipitate soluble in excess aqueous ammonia forming a deep blue solution.
  83. 83. • NB: the precipitate dissolves due to the formation of complex ions e.g. • Amphoteric hydroxides of Aluminiu. Zinc and lead dissolves in excess NaOH. • Pb(OH)2 (s) + 2OH- (aq) Pb(OH)4 2- (aq) + 2Na+ • Zn(OH)2 (s) + 4NH3(aq) {Zn (NH3)4}2+(aq) + 2OH-(aq) • Al(OH)3 (s) + OH- (aq) Al(OH)4 -(aq) + Na+ (aq) • Zinc and copper hydroxide dissolve in excess ammonia solution to form colourless tetra amine zincate and deep blue tetra amine copper respectively. • Zn(OH)2 (s) + 4NH3(aq) { Zn (NH3)4}2+(aq) + 2OH-(aq) • Tetramine zinc (II) ion. • Cu(OH)2 (s) + 4NH3(aq) { Cu(NH3) }2+(aq) + 2OH-(aq) Tetramine Copper (II) ion
  84. 84. How to make conclusions/inferences based on the observations • If on adding few drops of NaOH or ammonia solutions, white precipitate is formed, conclude by writing; • Either; Ca2+, Mg2+,Al3+, Zn2+,Pb2+ or Ba2+ ions present. • If green precipitate is formed, Fe2+ ions present. • If brown precipitate is formed, Fe3+ ions present. • If blue precipitate is formed, Cu2+ ions present.
  85. 85. •If on adding excess NaOH solution white precipitate persisted (insoluble in excess NaOH), conclude; Ca2+, Mg2+ or Ba2+ ions present. •If on adding excess NaOH solution white precipitate dissolves (soluble in excess NaOH) conclude. Al3+, Zn2+ or Pb2+ ions present. •If on adding excess ammonia solution white precipitate persisted (insoluble in excess), conclude; Either Ba2+, Ca2+, Mg2+, Al3+, or Pb2+ ions present. •If on adding excess ammonia solution white precipitate dissolves (soluble in excess) conclude; Zn2+ ions confirmed present.
  86. 86. For cations which exhibit similar properties in both dilute NaOH and aqueous ammonia. A confirmatory test has to be performed to distinguish them e.g. Al3+ ions and Pb2+ ions which form white precipitates which are soluble in excess NaOH but insoluble in excess aqueous ammonia. To distinguish them, two to three drops of either sodium chloride (NaCl), sodium sulphate (NaSO4) or potassium iodine KI are added. Lead ions form insoluble chloride, sulphate or iodide respectively.
  87. 87. Apart from using NaOH and ammonia solutions to identify cations flame test can also be used.
  88. 88. Addition of dilute acids Salt solutions reacts with dilute acids especially HCl & H2SO4, hence are used to confirm the preference of cations.
  89. 89. Test Observations Inferences 1.Add dilute HCl to solution - White ppt formed Pb2+,Hg2+ or Ag+ ions to be present - No white ppt is formed Pb2+, Mg2+ and Ag+ absent 1.Add dilute H2SO4 - White ppt formed Pb2+, Ca2+ or Ba2+ present - No white ppt formed Pb2+, Ca2+ and Ba2+ ions absent
  90. 90. Flame test for some cations This method is based on the colours the cations render a flame when heated. Procedure. •Dip a nichrome wire into concentrated HCl acid. •Heat the end that has been dipped in the acid using a flame until it cease to impart colour on the flame. This is done to clean the wire. •Dip the cleaned wire into the solution containing the test ion and then heat it in the hottest part of the flame. The test ion imparts its characteristics colour. NB: platinum wire, metallic spatula or glass rod can be used instead of the nichrome wire.
  91. 91. Cation Colour of the flame produced Lithium (Li+) Crimson (deep red) Sodium (Na+) Yellow Potassium (K+) Lilac (Purple) Calcium (Ca2+) Red Copper (Cu2+) Blue-green Barium (Ba2+) Apple green
  92. 92. CONFIRMATION TESTS FOR ANIONS • Anions are negatively charged ions • Most of the anions are radicals • They include:- CO3 2-, HCO3 -, SO4 2- & NO3 - • Examples of non-radical anions are S2- & Cl-, Br-,I-
  93. 93. Anions Test Observations SO4 2- Add barium chloride and mix followed by dilute nitric or HCl acid - A white precipitate insoluble in dilute acid SO3 2- Add Barium nitrate (Ba(NO2)2 or Barium Chloride (BaCl2) and mix followed by dilute HNO3 or HCl - A white precipitate soluble in dilute acid
  94. 94. CO3 2-, orHCO3 - , Heat the solid sample - K and Na carbonates do not decompose all other carbonates decomposes on heating Add dilute acid - Effervescence of CO2 which from white ppt in lime water is evolved. S2- Add dilute acid to solid Test with lead acetate - A gas is produced - Gas blackens Lead (II) acetate paper
  95. 95. Cl- Add AgNO3 or Pb (NO3)2 followed by dilute HNO3 - White precipitate insoluble in HNO3 (aq) - Ppt soluble in excess aqueous NH3 solution NO3 - Add a small amount of freshly prepared ferrous sulphate solution. Gently pour conc. H2SO4 down the side of the ppt tube such that the solutions do not mix - A brown ring may be formed at the interface/junctio n of the two solutions.
  96. 96. NB: preliminary and confirmatory Tests for Chlorides, Nitrates, Sulphates and Sulphites. • Confirmatory Tests for Nitrate Ions (NO3 -) • When the nitrates of sodium potassium are heated, the nitrates of the metals and oxygen gas are produced. • 2NaNO3 (s) 2NaNO2 (s) + O2 (g) • If the test solid is heated in a closed boiling tube, and a glowing splint is inserted at the mouth of the tube, the splint relights.
  97. 97. • Nitrates of heavy metals decompose on heating to produce brown nitrogen (IV) oxide gas, the respective metal oxide and oxygen. • 2Pb(NO3)2 2PbO (s) + 4NO2(g) + O2 (g) The Brown Ring Test • Freshly prepared solution of iron (II) sulphate to the test solution, followed by concentrated sulphuric (VI) acid. If a brown ring is formed at the interface of the two solutions, this indicates the presence of nitrate ions in the test solutions.
  98. 98. PRELIMINARY TESTS FOR ORGANIC COMPOUND • Organic compounds includes: • Hydrocarbons • Alkanols • Alkanals etc Flame test is used to test for hydrocarbons and alkanals (alcohols).
  99. 99. Blue non- sooty flame shows presence of c c Yellow smoky flame shows presence of c c or c c
  100. 100. TEST OBSERVATION INFERENCE PH of the aqueous solution pH lesser than 7 Weak carboxylic acid pH 4 to 6. O (R C O H) pH greater than 7 Compound is either an alkanol R-OH, ester O (R C O R) or an alkali metal salt of carboxylic acids O (R C ONa)
  101. 101. Burning Burns with a non luminous blue A saturated hydrocarbon with low carbon content C C Burns with a luminous yellow and sooty flame An unsaturated organic compound C =C or C C
  102. 102. Heating a mixture of the test substance with sodalime Colourless gas which is evolved burns with a luminous flame A salt of a carboxylic acid R-C-ONa The gas evolved is an alkane (saturated) R-C- or C-
  103. 103. Adding sodium carbonate to the test substance Colourless gas is evolved which forms a white precipitate with lime water Compound is a carboxylic acid or ester which is hydrolysed O O R-C-OH or R- C-O-R No reaction occurs Compound not a carboxylic acid or an ester
  104. 104. Adding acidified potassium manganate (VII) Decolourisation occurs in the cold Unsaturated compound or an alkanol ( areducing agent) c=c , c c or R OH Decolourisation occurs on warming Oxalates O O -O – C- C - O- or oxalic acid H2C2O4 Decolurisation occurs Unsaturated compound C =C or –C=C-
  105. 105. Adding carboxylic acid followed by sulphuric (VI) acid then warming the mixture Oily pleasant smelling compound formed The compound is an alkanol. R-OH
  106. 106. Note: •Potassium chromate (VI) could be used in all cases where potassium manganate (VII) is used. In all these cases, orange chromate (VI) ion changes to green chromate (III) ion. •Tests on organic compounds are based on the physical of the compounds: that is colour, smell, physical state at room temperature, and the solubility in water. The tests do not, therefore, give much indication regarding the identity of the compound being investigated; they give the functional regarding the identity of the compound being investigated; they give the functional groups series which is not a clear answer in most cases.
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The contents of this ppt are useful in revision of all the potential examination area during KCSE or any high school exam


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