Chemistry for Cambridge IGCSE Practical Lab Equipment PDF

Lab Equipment Advantages Disadvantages LAB EQUIPMENT Advan...

Lab Equipment Advantages Disadvantages LAB EQUIPMENT Advantages / Apparatus Name Use Disadvantages More accurate than a measuring Accurately measures cylinder volume of liquids to 25 Measures a fixed volume Pipette cm3 Slow / Solution slow to run out of pipette Accurately measures More accurate than a measuring Burette volume of liquids to cylinder 50cm3 Slower than measuring cylinder Measuring volumes of Measuring liquids and available in Easy to use / quick cylinder 25 cm3, 50 cm3, 100 cm3 Not accurate and 250 cm3 Holds and measures chemical liquid samples, Conical flask these chemicals can be Easy to mix heated, mixed and boiled in it. Used to store, mix, stir, Difficulty in swirling / mixing / Beaker pour, heating and shaking (compared to a conical transporting liquids flask) 1  Practical Lab Equipment Dropper Adds liquid in drops (Teat pipette) Gas jar For collecting gases It holds liquid (water) for collecting pure gas Trough samples over water To rinse various pieces of laboratory glassware Wash bottle OR Used to wash or rinse salts to remove impurities Grinding / crushing solid substances into smaller Mortar & pestle pieces (powder) Lid Melting or heating solid chemicals over a burner at a very high temperature Crucible Supporting and holding glassware at variable Stand with clamp heights Stand 2 Lab Equipment Round bottomed Heating and boiling liquids such as in distillation flask or other reagent reactions Heating and boiling liquids in distillation or other reagent reactions Flat bottomed flask (They are not as durable as round bottom flasks but have the sharp and vulnerable corners) Preparing a solution accurately to a known, Volumetric flask specific volume Gas syringe Collects & measures the volume of a gas Evaporating dish Used in heating solution to allow evaporation of (basin) chemical solutions Thermometer Measuring temperature of solid, liquid and gas For evaporation Watch glass Weighing solids Used as a lid for flasks and beakers Top pan balance Measuring mass of substances very accurately 3  Practical Lab Equipment Stopwatch Measuring time taken To seal the openings of test tubes, flasks and other laboratory glassware (insulation material) Stopper cork To prevent the gas escaping It is very light and porous To seal the openings of test tubes, flasks and other laboratory glassware Bung Used to prevent fluid passing through the neck of a test tube It is made of rubber Helps spread heat off the flame out for more even Gauze heating Used to support devices (beakers) which is placed Tripod above the Bunsen burner to heat and boil chemicals Produces flame for heating / Warming a Bunsen burner mixture Used for separating insoluble solid from a liquid Funnel with the help of a piece of filter paper 4 Lab Equipment For separating immiscible liquids of different Separating funnel densities in a reaction mixture Stirring rod / Glass Stirring / mixing chemicals rod Used for storing, mixing, and heating small Test tube amounts of chemicals Test tube holder Used to hold test tubes while heating To hold upright multiple test tubes at the same Test tube rack time Cools and condenses the vapour to be collected Condenser as a liquid To add small volumes of liquids to an exact Thistle funnel position 5  Practical Lab Equipment To hold a hot crucible, flask, evaporating dish, or a Metallic tong small beaker firmly Spatula To transfer small quantities of solid substances Knife To cut soft metals (group 1) Used for heating, combustion and sterilizing certain lab equipment Spirit burner Produces an open smaller flame than the Bunsen burner (Uses alcohol as a fuel) Used in the distillation of liquid mixtures to separate the mixture into its component parts, or Fractionating column fractions, based on the differences in volatilities (the more volatile move toward the top while the higher boiling stay towards the bottom) Used in measuring the pressure of liquids U-tube If water produced it will be produced as steam, so U-tube is used & sorrounded by ice to cool down the water (colorless liquid appear in U-tube). Allows a gas to pass from one container into another Delivery tube Suitable for many types of gas collection experiments To heat and boil small quantities of liquids over a Boiling tube flame Wider and thicker than test tube Filter tube Used to pass the gas through a liquid 6  Lab Equipment Tap funnel Used to separate the immiscible liquids Protects the eye area by preventing particles or Goggles dangerous materials from entering the eye while stirring, pouring or mixing IMPORTANT QUESTIONS 1. Explain one disadvantage of using a beaker instead of a conical flask. ✓ Difficulty in mixing /swirling, so conical flask is a better option when mixing /swirling. 2. Suggest and explain the effect on result when using burette instead of measuring cylinder? ✓ More accurate than measuring cylinder. 3. What is used to transfer any solid in any experiment? ✓ Spatula. 4. What is the purpose of watch glass? ✓ Prevents splashing of liquids and decrease evaporation. IMPORTANT NOTES Thermometer is used to measure the temperature of liquid and gas: ✓ To measure the temperature ✓ To measure the temperature of liquid: of gas: Should be placed in the liquid. Shouldn’t be placed in liquid. 7  Practical Separation Techniques SEPARATION TECHNIQUES Simple Fractional Separating Filtration Crystallization Distillation Distillation Funnel 1. Filtration ▪ Used to separate: 2 Solids (Soluble & Insoluble) E.g.: Mixture of sand and salt ▪ Ways to separate (steps): 1. Dissolve the mixture in excess to ensure complete dissolving. 2. Pour the mixture through filter funnel until all liquid drain. 3. Obtain solid on the filter paper as residue (insoluble). 4. The solution that passed through filter paper in conical flask is called filtrate (soluble). 8  Separation Techniques 2. Crystallization ▪ Used to obtain crystals: Soluble solid from water (As Crystals) Solid is obtained but the liquid is lost E.g.: Mixture of salt and water Solution Lamp ▪ Ways to obtain the crystals (steps): The solid is dissolved in a solvent The solution is heated to evaporate most of the solvent 1. Half evaporation (In evaporating dish) till point of crystallization. The cold solution is poured off to The hot solution is allowed to cool, obtain the crystals. The crystals 2. Leave to cool till crystals form. the solid appears as pure crystals may be dried by pressing them between sheets of filter paper 3. Filter to get crystals. 4. Dry between 2 filter papers gently. 3. Simple Distillation ▪ Used to separate: Soluble Solid from Liquid E.g.: Mixture of water and sodium chloride. ▪ Ways to separate (steps): 1. Boil the mixture 2. Condense the vapour through condenser and collect the liquid in a beaker. 3. The salt remains in the flask. HINT: You may observe color change. 9  Practical Separation Techniques 4. Fractional Distillation The idea depends on the fact of 2 or more liquids with different boiling point ▪ Used to separate: Two miscible liquids E.g.: Ethanol and water ▪ Ways to separate (steps): 1. Heat the mixture of both liquids. 2. Add both liquids to the flask. 3. The one with lowest boiling point evaporates first (as ethanol 78 degree) then condenses and is collected. 4. Then the liquid with higher boiling point (water at 100 degree) evaporates and condenses. NOTE: 1. If there is a flammable liquid (Ethanol), use electrical heater instead of Bunsen burner. (to prevent the risk of fire) 2. If water bath used instead of Bunsen burner you cannot obtain any liquid boils above100 degree as water boils at 100 degree. 10 Separation Techniques 5. Separating Funnel ▪ Used to separate: Two immiscible liquid E.g.: Mixture of water and oil ▪ Ways to separate (steps): 1. Let the mixture settle into 2 layers. 2. Open the tap to let only bottom liquid (denser one) run into beaker. 3. Close the tap then change the beaker. 4. Open the tap again to let upper layer run. IMPORTANT QUESTIONS: 1. How would you know when to stop heating during crystallization? ✓ When dipping a glass rod and observing the formation of crystals, or when crystals start appearing on the edge. 2. Why are crystals not evaporated to dryness? ✓ To avoid loss of water of crystallization. 3. What is meant by concentrated (little water) solution? ✓ More solute than solvent. 4. How can you make solution more concentrated? ✓ Evaporate to decrease amount of water (crystals are seen). 5. How could you show that solution is saturated? ✓ Add more solid, stirr, no more dissolves and solid is visible. 6. Explain the term “Decant” ✓ Pour off the liquid, leaving the solid. 11  Practical Separation Techniques CHROMATOGRAPHY ▪ Used to separate: 2 or more dissolved solids E.g.: - Colors - Amino acids - Sugars ▪ Ways to separate: 1. Get a Chromatography paper. 2. Draw a line with pencil called baseline (do not use ink as it will interfere with results). 3. Put a drop of the sample on the baseline using glass rod. 4. Put the paper in suitable solvent (level of solvent MUST be below baseline to avoid dissolving of the spot in solvent). 5. Leave the chromatogram for several hours. 6. Several spots will be separated. 12 Separation Techniques The Purity and identity of substances: - The distance moved by a particular spot is measured and related to the position of the solvent front. - The ratio of these distances is called retention factor (Rf) value. This value is used to identify the substance. 𝑫𝒊𝒔𝒕𝒂𝒏𝒄𝒆 𝒎𝒐𝒗𝒆𝒅 𝒃𝒚 𝒕𝒉𝒆 𝒔𝒖𝒃𝒔𝒕𝒂𝒏𝒄𝒆 Rf= 𝑫𝒊𝒔𝒕𝒂𝒏𝒄𝒆 𝒎𝒐𝒗𝒆𝒅 𝒃𝒚 𝒕𝒉𝒆 𝒔𝒐𝒍𝒗𝒆𝒏𝒕 𝒇𝒓𝒐𝒏𝒕 - Paper chromatography is one test that can be used to check the purity of a substance. If the sample is pure, it should give one spot by chromatography. - The identity of substances can also be checked by comparing its Rf value to that of a sample we know to be pure. - The most common tests for purity are the melting and boiling points. Impurities would lower the melting point and raise the boiling point of a substance. - The process of purification is very important in many areas of chemical industry such as medical drugs. 13  Practical Separation Techniques IMPORTANT NOTES: Substances separate according to their solubility in the solvent. E.g.: The dyes are carried by the solvent and begin to separate. The substances that is most soluble moves fastest up the paper. The run stops just before the solvent front (the level reached by the solvent). An insoluble substance will remain at the origin. If sample is insoluble in water, USE ethanol (organic solvent). Make sure to cover with lid if using ethanol as ethanol is volatile (can evaporate easily). If only 1 spot produced, then the sample is pure. If the separated substances are colorless, the paper is treated with locating agent that reacts with the sample to produce coloured spots. E.g.: Amino acids are colorless, a locating agent is used NOTE: Molecules of amino acids can also be viewed under ultra violet. IMPORTANT QUESTIONS 1. Why is a lid necessary on top of the beaker, in case of using ethanol as a solvent? ✓ To prevent evaporation / loss of solvent. 2. If the dyes are insoluble in water, name a suitable solvent that could be used. ✓ Organic solvent (E.g.: Ethanol) 14  Summary Still confused?? How to decide which method to use?! Here is a key to help 15  Practical Electrolysis ELECTROLYSIS It is a breaking down of ionic compound when molten or aqueous solution using electricity. Diagrams to consider: oxygen gas O2(g) hydrogen gas H2(g) 1 volume of gas 2 volumes of gas Acidified water (dilute sulfuric acid) anode cathode (graphite) (graphite) Compare volumes of O2 & H2: Volume of H2 is double volume of O2 Electrolysis of aqueous solutions: 1. Concentrated: Ions discharged: Less reactive (at cathode) & Less complex (at anode). (Revise Reactivity well) 2. Diluted: Water is electrolyzed (broken down). Ions discharged: H+ & OH-. 16  Electrolysis Observations to consider: If bulb used in circuit ➔ It will light. If gas produced ➔ Bubbles / Fizz / Effervescence of (color) gas will be seen. E.g.: If chlorine ➔ Bubbles of green gas. If metal formed at electrode ➔ color of metal (solid) formed will appear. E.g.: If copper ➔ Red-brown metal is formed NOTE: During electrolysis of concentrated aqueous copper sulfate, color changes from blue to colorless. IMPORTANT QUESTIONS 1. Name the process of breaking down a substance using electricity ✓ Electrolysis 2. Which substance breaks down during electrolysis of dilute solutions to form O2 and H2 gases? ✓ Water 3. In the electrolysis of molten zinc chloride using inert electrodes, suggest why iron electrodes cannot be used in this experiment. ✓ Iron reacts (with chlorine) / iron is not inert IMPORTANT NOTES The change in electrolyte of electrolysis of dilute acid ➔ PH decreases, becomes more acidic (more concentrated) as water is electrolyzed. 17  Practical Electrolysis ELECTROPLATING Steps to electroplate a small metal key with silver: - Clean the metal key using sand paper. - Key in negative pole (cathode). - Anode is silver metal. - Electrolyte is silver salt as Silver Nitrate. - Switch on electricity. - Rotate the key. - Dry. WHY DO WE? 1. Clean using sand paper? ✓ To let the metal layer stick firmly. 2. Rotate the object? ✓ To let the metal layer forms evenly on all sides. 3. Control current and temperature of electrolyte? ✓ To prevent formation of the metal layer too fast ,so the it doesn’t flake off. 18  Acid and Bases ACID AND BASES Acid: a substance that dissolves in water to produce hydrogen ions (proton donor). This solution: Contains an excess of H+ ions. Turns litmus paper red. Has a pH lower than 7. Base: a substance that dissolves in water to produce hydroxide ions (proton acceptor). This solution: Contains an excess of OH- ions. Turns litmus paper blue. Has a pH higher than 7. Comparison between Strength of acids Method Strong/conc. Weak/dilute Universal indicator Red Yellow PH meter Low 1-3 High 4-6 Metal Carbonate (MgCO3) rapid effervescence (CO2) slow effervescence (CO2) Electrical conductivity High Low Add Alkali Higher temperature rise is more in conc/strong Comparison between Strength of Base Method Strong/conc. Weak/dilute Universal indicator Violet Dark green PH meter High 11-14 Low 8-10 Ammonium Salt (NH4Cl) rapid Bubbles(NH3) slow Bubbles (NH3) Electrical conductivity High Low Add Acid Higher temperature rise is more in conc/strong 19  Practical Acid and Bases You can measure the acidity of a substance by using: 1. INDICATORS Color in each state Name of indicator Acid Neutral Base Litmus paper Red Violet Blue Methyl orange Red Orange Yellow Phenolphthalein Colorless Colorless Pink 2. UNIVERSAL INDICATORS 1,2,3 4,5 5-6 7 8-9 10-11 12,13,14 Weak Very weak Very weak Weak Strong Strong acids Neutral acids acids alkali alkali alkalis Red Orange Yellow Green Dark green Blue Violet IMPORTANT QUESTIONS 1. Suggest why Universal Indicator is not a suitable indicator. ✓ No sharp colour change / No (clear) end point 20 Reaction of Metals REACTION OF METALS 1. Reactions with acids a) Metal (more reactive metals) + Acid ➔ Salt + Hydrogen Observation: Bubbles of H2 Example: 2 Na + 2 HCl ➔ 2 NaCl + H2 b) Metal carbonate + Acid ➔ Salt + Water + Carbon dioxide Observation: Bubbles of CO2 Example: CaCO3 + 2 HCl ➔ CaCl2+CO2+H2O c) Metal oxide / Hydroxide (base) + Acid ➔ Salt + Water Example: NaOH + HCl ➔ NaCl + H2O CaO + 2HCl ➔ CaCl + H2O 2. Reactions with water a) Metal (most reactive) + Cold water ➔ Metal hydroxide + Hydrogen Observation: Bubbles of H2 Example: 2 K + 2 H2O ➔ 2 KOH + H2 2 Na + 2 H2O ➔ 2 NaOH + H2 Ca + 2 H2O ➔ Ca(OH)2 + H2 NOTE: ▪ K, Na react vigorously with cold water ▪ Ca reacts readily but not violently with cold water ▪ Mg reacts with cold water slowly (to form Magnesium Hydroxide and Hydrogen) ▪ Mg reacts faster with steam producing Magnesium Oxide and Hydrogen b) Metal (Zinc & Iron) + Steam ➔ Metal oxide + Hydrogen Example: Zn + H2O ➔ ZnO+H2 Observation: Bubbles of H2 21  Practical Reaction of Metals 3. Reaction with oxygen a) Metal + Oxygen ➔ Metal oxide Example: 2 Cu + O2 ➔ 2 CuO Observation: Black solid / ash formed Example: 2 Mg + O2 ➔ 2 MgO Observation: White solid / ash formed 4. Reducing copper oxide a) Reducing Copper (II) oxide to copper CuO + H2 ➔ Cu + H2O (Black) (Red brown) Observation: ▪ A red-brown solid is formed. ▪ A colorless liquid is formed in the U- tube (water) What is the purpose of the ice surrounding the U-tube? ✓ Acts as a condenser to cool down steam and a colorless liquid (water) will appear in the U-tube. REACTIONS OF GROUP 1 METAL: Lithium Sodium Potassium All three are less dense than Floats Floats Floats water Fizzes Fizzes Fizzes Hydrogen gas is produced Keeps its shape as Melts into a ball Melts into a ball as The reactions are all it reacts and gets and catches fire it reacts exothermic smaller with a lilac flame increasing vigour of reaction Least reactive Most reactive of the three All three metals disappear as they react, forming soluble hydroxides as products 22  Reaction of Metals IMPORTANT NOTES Only metals above hydrogen in the reactivity series can react with acids. It is very dangerous to react K / Na with acid or steam as that would cause explosion Order of reactivity from strongest: 1. Explosively 2. Violently 3. Vigorously IMPORTANT QUESTIONS 1. Suggest why sodium and potassium are not used in experiment? ✓ Too reactive / dangerous / vigorous reaction 2. Why we use Cotton / Mineral Wool? ✓ To allow escaping of gas. ✓ To hold & absorb water. ✓ To prevent splashing. 3. Why mixture is always stirred with glass rod instead of spatula? ✓ Glass is unreactive as if metal is used it may react with spatula. 4. Suggest and explain one improvement to this experiment (in figure 1) ✓ Use cotton thread to hold a test-tube containing the acid in the flask; to prevent the collection of air 5. Suggest why the reading on the measuring cylinder was 30cm3 after the acid had been added and before the timer had been started. (in figure 1) Figure 1 ✓ Air is displaced when the acid is added. 23  Practical Rate of chemical reactions RATE OF CHEMICAL REACTIONS Introduction ▪ The reaction starts fast as the concentration / Mass of reactants is high (Curve Steepest) ▪ The reaction then slows down as the concentration / Mass is decreasing (Curve less Steep) ▪ The reaction is finished when limiting factor is used up (all reactants react) (Curve Flat) Factors affecting the rate of reactions Temperature Pressure Concentration Surface Area Stirring Catalyst Light By increasing any of these factors the rate of chemical reaction increases 24 Rate of chemical reactions Different reactions with different rate Graph A Steeper curve ➔ Higher rate of reaction. Higher end point ➔ Higher concentration / Mass of limiting factor. Graph B Steeper curve ➔ Higher rate of reaction ➔ Using higher temprature / adding catalyst / increasing concentration / using smaller lumps Graph C Less steep curve ➔ Slower rate of reaction ➔ Using lower temprature / decreasing concentration / using larger lumps Same endpoint ➔ Same concentration / Mass of limiting factor 25  Practical Rate of chemical reactions Anomalous Point: This may happen due to misread the measuring cylinder or read too late If the point is below the curve, then it happened due to taking late reading This is the way to obtain the right measurement Some graphs BEST FIT LINE SMOOTH LINE CURVE Draw it with pencil using Ruler Draw it with pencil without using Ruler It is not necessary to path by all points 26 Rate of chemical reactions BAR CHART Draw it with pencil using Ruler There are gaps between columns The difference between Bar Chart and Histogram: 27  Practical Rate of chemical reactions IMPORTANT NOTE WHY Polystyrene is better than copper can? Polystyrene is an insulator which will decrease heat loss and therefore the temperature measure high. IMPORTANT QUESTIONS 1. Suggest an advantage of taking readings regularly at short time intervals (e.g. every 15 seconds) ✓ There will be more readings so a smoother curve and more accurate graph will be drawn. 2. Suggest an advantage and disadvantage of taking readings after 1 min? ✓ Advantage: fair experiment. ✓ Disadvantage: may be the reaction not finished as if you are measuring temperature, termperature may be still changing. 3. Why does the rate of the reaction decrease over time? ✓ Concentration of reactants decrease 4. When does the reaction finish? / Explain why the rate is zero? ✓ All limiting factors react / used up / decompose 5. Why there is an anomalous point? ✓ It can be a result of misreading measuring cylinder or reading too early / late. 28 Rate of chemical reactions 6. What measurements should be taken to follow the rate of the reaction? ✓ Decreasing of volume / mass of reactants (E.g. Hydrogen peroxide and acid) or Increasing of volume / mass of products (E.g. Gas produced) ✓ Time taken 7. How could a student separate the catalyst from the reaction mixture at the end of the reaction? ✓ Filtration 8. Why we crush the sample of solid before any experiment? ✓ To increase the surface area so the rate increases. 9. What type of chemical reaction occurs when magnesium ribbon reacts with dilute hydrochloric acid? And what is the observation? ✓ The reaction is Exothermic / Redox / Displacement ✓ Observations Rapid effervescence and the test-tube felt hot 10. How could you know than any reactant is in excess? ✓ Solid is visible 11. The volume of gas collected is less than expected, Give reasons for that? ✓ Volume of gas less / lower ✓ Gas/carbon dioxide dissolved in water 12. What is the reason of cracking the tube while heating the magnesium ribbon? ✓ Releasing of large amount of energy. 13. How the dissolving process could be speeded up? ✓ By stirring or heating. 14. Explain why copper can is used instead of a boiling tube? (asked when there is a temperature measurement) ✓ Copper is a good conductor of heat as it is a metal, so higher temperature change. 29  Practical Preparation of Salts PREPARATION OF SALTS To prepare any salt you should consider: Parent Acid & Soluble Salt Titration (Group 1) Alkali or Basensoluble Parent Acid & Nuetralization Soluble Salt Alkali or Basensoluble 2 Soluble Preceptation Insoluble Salt Salts Solubility Rules: Soluble Insoluble All Group 1 and ammonium salts All Nitrates Silver Chlorides Chlorides Lead Chlorides Calcium Sulphates Sulphates Barium Sulphates Lead Sulphates Sodium and potassium (Group 1) Carbonates Ammonium Carbonates Sodium, potassium, magnesium and Oxides calcium oxide NOTE: As temperature decrease, solubility decrease 30  Preparation of Salts Methods of Preparations: 1. Neutralization ▪ Stage 1: An excess (more than enough) of the solid is added to the acid and allowed to react. Using an excess of the solid makes sure that all the acid reacts. ▪ Stage 2: The excess solid is removed by filtration. ▪ Stage 3: The filtrate is gently evaporated to concentrate the salt solution. ▪ Stage 4: When crystals can be seen forming (crystallization point), heating is stopped and the solution is left to cool. ▪ Stage 5: Wash the crystals and dry between two filter papers or in oven. CuCO3 ✓ The difference between using Metal Carbonate and Metal Oxide in neutralization: Metal Carbonate Metal Oxide Reaction finished No more bubbles No more Solid dissolved Observation Bubbles - No heat is needed Heating to increase the rate Heat (as reaction is fast at room of the reaction temperature) 31  Practical Preparation of Salts 2. Titration ▪ Stage 1: The acid solution is poured into a burette. A known volume of alkali solution is placed in a conical flask using a pipette. A few drops of an indicator are added to the flask. ▪ Stage 2: The acid solution is run into the flask from the burette until the indicator just changes color. Notice the end-point for the reaction, the volume of acid run into the flask is noted. ▪ Stage 3: The experiment is then repeated without using the indicator. - The same known volume of alkali is used in the flask. - The same known volume of acid as noted in the first parts is then run into the flask. ▪ Stage 4: The salt solution is evaporated and cooled to form crystals. ✓ Titration can be used to compare the strength of acids and bases: Example 1: If we have same volume of acids (acid A and acid B) with different types Base Base 4 cm3 of base used 8 cm3 of base used Acid A Acid B Experiment using acid A Experiment using acid B Observation Acid A took 8 cm3 of base to reach neutralization. Acid B took 4 cm3 of base to reach neutralization. Conclusion Acid A is twice stronger than acid B. NOTE: For fair comparison use same concentration of base and same starting point. This experiment is also used to compare the concentration of acids. 32  Preparation of Salts Example 2: If we have different types of bases (base X and base Y) Base X 2 cm3 of base used Base Y 4 cm3 of base used Acid Acid Experiment using base X Experiment using base Y Observation 2 cm3 of base X enough to neutralize acid. 4 cm3 of base Y enough to neutralize acid. Conclusion Base X is stronger than base Y. NOTE: For fair comparison use same volume and concentration of acid. This experiment is also used to compare the concentration of bases. 3. Precipitation Soluble Soluble Insoluble Soluble Salt Salt Salt Salt ▪ Stage 1: Add two salts together (Corresponding sodium salt and corresponding nitrate). ▪ Stage 2: Stir till you observe precipitate forms at the bottom of the beaker. ▪ Stage 3: ❖ To obtain the residue: Filter the mixture to obtain insoluble salt (residue). Then wash the residue and dry between two filter papers or in warm oven. If you want to obtain soluble salt (the filtrate) → Wash the residue to collect all the traces remaining of the filtrate, then evaporate till crystallization point, then the crystals are left to cool and filtered then washed and dried between two filter papers. 33  Practical Preparation of Salts IMPORTANT QUESTIONS 1. Why we rinse the burette, pipette, measuring cylinder, etc… between each experiment? ✓ To remove traces of impurties (such as acid and base) 2. What would be the effect of adding water in burette in titration? ✓ The volume used will be increased because the concentration is decreased. 3. What would be the effect of adding drops of water in conical flasks in titration? ✓ Doesn’t affect the value because it is the same amount (n). 4. What would be the effect on the results of warming solutions (acid) in a reaction before titration? ✓ No effect, as there is no change in concentration, only concentration of the limiting factor affects the results. 5. How did the student know that all of the acid had reacted in neutralization? ✓ Solid remain / metal oxide stops dissolving / reacting. 6. The titration between sodium hydroxide and hydrochloric acid, what would be the effect of rinsing the conical flask with NaOH? ✓ The volume of acid needed increases as the volume of NaOH increased. 7. The titration between sodium hydroxide and hydrochloric acid, what would be the effect of rinsing the conical flask with water? ✓ No effect , amount doesn’t change. 8. How is the excess metal oxide removed in neutralization? ✓ Filtration / filter 9. Suggest how the method would differ if zinc carbonate were used instead of zinc oxide to make zinc chloride crystals by reacting with dilute hydrochloric acid. ✓ Heating / warming not necessary. 34  Colors COLORS FOR ELEMENTS Metals: All metals are SILVERY grey except gold (Yellow) and copper (Red brown). Non-metals: State of matter at Non-metals Color room temperature Carbon Solid Black Sulphur Solid Yellow Boron Solid White Phosphorus Solid Red Yellow Astatine Solid Black Iodine Solid Grey Bromine Liquid Reddish brown Chlorine Gas Green Fluorine Gas Yellow Nobel gases Gas Colorless Hydrogen Gas Colorless Oxygen Gas Colorless Nitrogen Gas Colorless Iodine sublimes giving violet gas and produces brown color in solution. 35  Practical Colors COLORS OF COMPOUNDS 1. Covalent compounds: They are colorless except nitrogen dioxide NO2 (Brown fumes). 2. Ionic Compounds: Compound Color Compounds of group 1,2,3 White solid and colorless on solution Cupper II oxide CuO Black Cupper II hydroxide Cu(OH)2 Blue Anhydrous cupper II sulphate CuSO4 White Hydrated cupper II sulphate CuSO4.5H2O Blue crystals Iron II hydroxide Fe(OH)2 Green Iron III hydroxide Fe(OH)3 Reddish brown Potassium dichromate (oxidizing agent) Orange K2Cr2O7 Potassium permanganate (oxidizing agent) Violet KMnO4 All Cu II salts are blue in solution except CuCO3 green and CuO black. All Fe II salts are green in solution All Fe III salts are reddish brown in solution Transition Metals have colored compound How can you distinguish between alkenes and alkanes? Add bromine water: Alkenes ➔ Turns colourless Alkanes ➔ No change 36  Anions and Cations ANIONS AND CATIONS Test for Anions Anion Test Results Carbonate Effervescence as carbon Add dilute acid CO32- dioxide produced Chloride White Cl- Bromide Acidify with dilute nitric acid then Cream Br - add aqueous sliver nitrate Iodide Yellow I- Nitrate Add aqueous sodium hydroxide, then aluminium foil, warm Ammonia produced NO3- carefully Sulfate Acidify, then add aqueous barium White SO42- nitrate Sulfur dioxide produced will Sulfite Add dilute hydrochloric acid, turn acidified aqueous warm gently and test for the SO32- potassium manganate (VII) presence of sulfur dioxide from purple to colorless 37  Practical Anions and Cations Test for Cations Effect of aqueous sodium Effect of aqueous Cation hydroxide ammonia White ppt., Aluminium White ppt., Soluble in excess giving a Al3+ Insoluble in excess colorless solution Ammonium Ammonia produced on warming - NH4+ Calcium White ppt., No ppt or very slightly White Ca2+ Insoluble in excess ppt. Chromium (III) Green ppt., Green (grey green) ppt., Cr3+ Soluble in excess Insoluble in excess Light Blue ppt., Copper Light Blue ppt., Soluble in excess giving a dark Cu2+ Insoluble in excess blue solution Iron (II) Green ppt., Green ppt., Fe2+ Insoluble in excess Insoluble in excess Iron (III) Red-Brown ppt., Red-Brown ppt., Fe3+ Insoluble in excess Insoluble in excess White ppt., White ppt., Zinc Soluble in excess giving a Soluble in excess giving a Zn2+ colorless solution colorless solution 38  Flame Test FLAME TEST Steps for flame test preparation: 1. Get a nickel-chromium or platinum wire. 2. Dip the wire in hydrochloric acid / nitric acid. (To make sure it is clean) 3. Dip the clean wire in powder or solution of the ionic metal salt then introduce it to the bunsen burner. 4. The color of flame is observed. Flame Test Results: Lithium Sodium Potassium Calcium Barium Copper Metal ion (Li+) (Na+) (K+) (Ca2+) (Ba2+) (Cu2+) Orange- Light- Blue- Flame color Red Yellow Lilac Red green green Diagram Sources of error of flame test: ▪ Impurities affect the result. ▪ Brightness of the flame various from one sample to another. 39  Practical Flame Test IMPORTANT QUESTIONS 1. In test and observation table, when you’re told that upon heating condensation is formed at the top of the tube (colouless drops /water formed at top of tube). ✓ Your conclusion should be that the solid is hydrated or contains water. 2. What will you observe when alcohols are touched with lighted splint? ✓ Catches fire and burns with yellow flame. 3. A Bunsen burner was used to heat solid. Describe how a Bunsen burner is adjusted to give a very hot flame. ✓ Open (air hole/collar). The Bunsen burner has three standard flames which depend on the position of the collar and use: Position of collar Flame Used for This is the safety flame Shimmering and bright Closed and is generally not yellow flame used Gentle heating E.g. solutions in a boiling Half-closed Steady blue flame tube which is moved into and out of flame Hottest flame, used in thermal Noisy, roaring blue Fully open decomposition flame reactions, oxidation, dehydration 40  Test for gases TEST FOR GASES Gas Test Result / observation Hydrogen H2 Lighted splint Pops Oxygen O2 Glowing splint Relights Carbon dioxide Lime water (Ca(OH)2 solution) Turns milky CO2 Chlorine Cl2 Litmus paper Bleaches / Turns white Turns blue Ammonia NH3 Red litmus paper Turns pH > 7 It has a pungent smell Acidified aqueous potassium Sulfur dioxide SO2 Turns from purple to colorless manganate VII (KMnO4) IMPORTANT QUESTIONS 1. Why shouldn’t we use hydrogen in excess? ✓ As hydrogen is flammable, It shouldn’t be used in excess. METHODS OF DRYING GASES 1. Concentrated Sulphuric acid used to dry all gases except Ammonia (base). 2. Anhydrous calcium chloride is used for all gases except Ammonia, which forms a complex with calcium chloride. 3. Calcium oxide is used to dry ammonia and neutral gases. 41  Practical Test for gases IMPORTANT QUESTIONS 1. Suggest one reason why the gas produced (gas in) is passed through concentrated sulfuric acid. ✓ To dry the gas / remove water and impurities. 2. Why is concentrated sulfuric acid not used to dry ammonia gas? ✓ Ammonia is base and will react with acid (neutralization). METHODS OF COLLECTING GASES: Downward Gas syringe Method Upward delivery Over water delivery Use Gas more dense Gas less dense Gas is sparingly To measure when… than air than air soluble in water volume Apparatus Carbon dioxide, Chlorine, Carbon dioxide, Ammonia, Any gas Examples Sulphur dioxide, Hydrogen, Hydrogen Hydrogen Oxygen Chloride 42  Precautions EXPERIMENTAL PRECAUTIONS If using / substance Hazard Safety precaution Strong acid Skin corrosion Wear gloves / goggles Strong alkali Skin corrosion Wear gloves / goggles -Keep away from flame -Well ventilated lab Alcohol, hydrocarbon, Flammable (catch fire) -Heat carefully hydrogen -Use electric heater or water bath Oxidant and provide Potassium manganate, oxygen which causes Keep away from potassium dichromate flammable material to flammable materials burn more fiercely Toxic gases Use fume cupboard E.g.: CO, SO2, NO, NO2, F2, Toxic / poisonous Well ventilated lab Cl2, Br2, I2 If heating a liquid / Spitting which could harm Heat carefully solution your eyes 43  Practical Test for cation ana anions Example 1 Adding water then dissolving occurs it means salt is soluble May be Group I, II, III or NH4+ Neutral, not acid nor base Test for halides Bromide ion Br- Sodium ion Na+ Sodium Bromide Transition Coloured Chromium Green metal Compound Green ion Cr3+ Test for cations (Chromium ion) Note: with excess aqueous sodium hydroxide, chromium ion is soluble but it is not an Green precipitate observation with excess green solution / dissolves 44 Test for cation ana anions Test for cations (Chromium ion) Note: with excess ammonia, chromium ion is insoluble (not an observation) which means it remains green precipitate Grey-Green precipitate Test for Anions (Chloride ion) White precipitate Fume cupboard / gloves / goggles 45  Practical Test for cation ana anions Example 2: (Try to solve it) 46 Test for cation ana anions 47  Practical Sources and Improvements SOURCES OF ERROR AND THEIR IMPROVEMENTS If using Improvement and explanation Add acid to water not vice versa as this Acid and Water exothermic reaction may cause beaker to be cracked. Metal ribbon or adding one spatula of Measure same mass of the substance a solid to every experiment using balance to be more accurate. Glass beaker to determine the temperature rise in exothermic Change glass beaker with polystyrene reactions such as neutralization cup to decrease heat loss. reaction Glass beaker to heat liquid using flame Change glass beaker with copper can Or (better heat conductor) to transfer Glass beaker contain water to more heat, so it will give more determine temperature rise caused by temperature rise. a fuel using spirit burner Rinse the burette with distilled water after adding the 1st liquid to remove A burette to add more than one liquid any traces of it, then rinse with the / solution separately after each other second liquid to remove any traces of water. Volatile liquid (ethanol) in Cover beaker with lid to prevent chromatography evaporation of solvent 48 Sources and Improvements Improvements to increase accuracy: Repeat the whole experiment and take average. Measure using burette / pipette instead of measuring cylinder. Decrease intervals time to take more readings. Use a digital thermometer. Use cotton thread to hold a test tube containing acid in the flask to make sure no air is collected. Use a lid to reduce heat. IMPORTANT QUESTIONS ✓ Why we rinse the burette, pipette, measuring cylinder, etc… between each experiment? ✓ To remove traces of acid ✓ Suggest how the reliability of the results could be checked. ✓ Repeat the experiment ✓ Compare or to check for anomalous results ✓ Suggest why experiments in which chlorine gas is prepared are done in a fume cupboard. ✓ Chlorine/gas is poisonous / toxic. ✓ A bottle of zinc chloride is labelled corrosive. State one safety precaution that should be taken when using zinc chloride. ✓ Wear gloves / goggles ✓ Why is it better to use an electric heater instead of a Bunsen burner to heat (E.g. ethanol)? ✓ As ethanol is flammable 49  Practical Sources and Improvements DANGER OF “SUCKING BACK”: There is danger of “sucking back” in the gas collected over water. The problem arises when heating is stopped before delivery tube is removed from water. i. The reduced pressure in the reaction tube as it cools. ii. The cold water can be sucked back into the hot boiling tube, the tube will crack and an explosion may occur. “Sucking Back” Can be prevented by making sure that the delivery tube is removed first before heating is stopped. IMPORTANT QUESTIONS 1. Why must the delivery tube be removed from the water when the heating is stopped? ✓ To prevent suck back (of water) 2. What is the purpose of the suction pump? ✓ To suck gases/ products through apparatus, to prevent explosion / cracking of test tube. 50  Investigation INVESTIGATION These ideas are intended to help you grasp and understand the basic steps to follow in different experiments 1. Rusting Questions often want you to compare the effect of different types of water on rusting). Our target here is to measure the change in mass and build conclusions accordingly. - weigh specified number of nails (to get initial mass) - place them in a suitable container (test-tube) and add water - Leave for 1 week - dry the nails - reweigh the nails (to get final mass) - repeat with the other type of water - greatest increase in mass means highest rusting has occurred NOTE: We use same steps for comparing the effectiveness of protection methods (e.g. painting) but we repeat using the other type of protection instead of water 51  Practical Investigation 2. Reactivity of metals Questions often want you to compare the order of reactivity of metals. Our target here is to measure the volume of gas produced in a fixed time (rate) and draw a proper conclusion. - use fixed mass of metal - use fixed volume of acid - add metal to acid in a suitable container (conical flask) - measure volume of gas produced in a specific time - using gas syringe and stop clock - repeat with the other metals - biggest volume of gas produced in a specific time (highest rate) is most reactive NOTE: Rate = Volume of gas / time Alternatively, you can measure the change in temperature; the greatest change in temperature shows the highest reactivity. 3. Electroplating Questions ask you to coat a metal object with another metal e.g. coating a key with silver: - clean the metal object (key) using sandpaper - place the key in the negative pole (cathode) Electrolyte should be a - Silver rod connected to the positive pole (anode) solution of a soluble - electrolyte is silver nitrate compound containing the metal we want to coat - switch on electricity with (e.g. silver nitrate) - rotate the key tip: use nitrates as all nitrates are soluble - remove the key, wash and dry it 52  Investigation 4. Exo & Endo Questions often want you to determine whether reactions are endothermic or exothermic and which one has bigger energy change. Our target here is to find how temperature changes during the reaction and build conclusions accordingly. - measure initial temperature using a thermometer - add the other reactant - measure final temperature - if temperature increases then the reaction is exothermic - if temperature decreases then the reaction is endothermic - bigger temperature change is bigger energy change IMPORTANT QUESTIONS 1. During the reactions of different metals with an acid, which reaction produces largest increase in temperature? ✓ The most reactive metal produces largest increase in temperature 2. During the reactions of different acids with a metal, which reaction produces largest increase in temperature? ✓ The most concentrated / strongest acid produces largest increase in temperature 3. In reactions involving measuring change in temperature, why is polystyrene cup used not a copper can? ✓ As polystyrene is an insulator used to reduce heat loss, but copper is a good conductor NOTE: After the reaction has stopped, the temperature returns back to room temperature 53  Practical Investigation 5. Mixture of solids with different solubility in water Question mentions that solid X is soluble in water while solid Y is insoluble in water and asks you to calculate the percentage by mass of X & Y. - weigh a known mass of the mixture of solids - add water to the mixture in a suitable container Beaker (beaker) - stir - filter to get the residue (insoluble solid Y) Filter paper - wash and dry the residue Funnel - weigh the residue - percentage by mass of Y = (mass of residue/mass of Conical flask mixture) *100 - percentage by mass of X = (mass of mixture mass of residue/mass of mixture) *100 6. Salt preparation a) Precipitation To prepare an insoluble salt we use ‘precipitation’. You should think of two soluble salts where one salt contains the cation, and the other contains the anion e.g. if you want to prepare AgCl , we search for two soluble salts; one contains Ag+ & the other contains Cl-.We can thus use AgNO3 & NaCl. *for a soluble salt containing the cation X, use X NO3 (as all nitrates are soluble) for a soluble salt containing the anion Y, use NaY (as all sodium compounds are soluble) NOTE: We crush the sample before any experiment to increase the surface area. - mix equal volumes of the two soluble salts - filter to get the residue (insoluble salt) - wash the residue with distilled water - dry the residue between two filter papers or in warm oven 54  Investigation b) Neutralization To prepare a soluble salt (not from group 1) we use ‘neutralization’. We add an insoluble base (metal oxide/carbonate) to the acid. The required salt is soluble so we have to carry out crystallization to obtain crystals. Making the salt - add excess chosen insoluble base to a certain volume of chosen acid (if metal oxide add till no more solid dissolves and if metal carbonate add till no more effervescence) - warm and stir - filter to remove unreacted insoluble base (residue) Obtaining crystals - half evaporation (of filtrate) till crystallization point - leave to cool down and form crystals - filter to get the crystals - wash with distilled water - dry between to filter papers or in a warm oven c) Titration To prepare a group 1 salt we use ‘titration’. The required salt is soluble so we have to carry out crystallization to obtain crystals. Making the salt - using a pipette add 25 cm3 of alkali to a conical flask - add few drops of an indicator (e.g. methyl orange) to the alkali - fill the burette with the acid - add the acid from the burette slowly till the indicator changes color (till reaching the endpoint) - record the volume of acid used from the burette - repeat the experiment without indicator using same volume of acid and alkali 55  Practical Investigation Obtaining crystals - half evaporation till crystallization point - leave to cool down and form crystals - filter to get the crystals - wash with distilled water - dry between to filter papers or in a warm oven 7. Titration for comparison For questions asking you to find the more concentrated *We use titration* e.g.comparing two cleaners A & B for which contains most concentrated alkali - measure a specific volume of cleanser A using a pipette - add cleanser A in a conical flask - add methyl orange in the conical flask - fill a burette with hydrochloric acid - add hydrochloric acid slowly from the burette till indicator changes colour - repeat with cleanser B using same volume - biggest volume of acid is most concentrated NOTE: for questions asking you to find: the concetration, the strong and weak, the one that contains more ions, the one that neutralizes more ,we use same steps 56  Investigation 8. Catalyst a) To prove that the catylst does not change during the reaction: - weigh a known mass of the catalyst before the reaction - allow the reaction to take place - filter, wash and dry the catalyst (if the reaction involves liquids or solutions) - reweigh the catalyst after the reaction b) To prove whether the catalyst speeds up the rate: e.g. an experiment that produces a gas - carry out the experiment without catalyst - measure the volume of gas produced in a specific time (rate) - use stop clock and gas syringe - repeat the experiment with catalyst - measure volume of gas produced in the same specific time (rate) - biggest volume of gas procuced in the specific time should be produced in the presence of catalyst 9. Obtaining water from a hydrated salt and calculating its percentage Do not forget to draw a diagram if required a. To collect the water and then calculate its percentage - weigh a known mass of the hydrated salt - heat the salt - allow condesation to take place - continue until no more water is collected - weigh water - percentage of water= (mass of water/ mass of hydrated salt)*100 57  Practical Investigation b. To calculate the percentage of water directly without collecting it - weigh a known mass of the hydrated salt - heat to remove water of crystallization - reheat till constant mass (to make sure all water evaporated) - weigh the remaining solid - percentage of water= (mass of hydrated salt – mass of remaing solid / mass of hydrated salt)*100 10. Fuels Questions ask you to compare which fuel/ coal produces more energy. Our target here is to burn the fuel, measure the temperature change that the fuel combustion results in , and build conclusions accordingly. - add the first fuel to spirit burner - measure the initial temperature of water - burn the fuel for a specific time - measure the final temperature of water - repeat with the second fuel Use same: *volume of fuel *volume of water *type of thermometer *distance between water and fuel/coal - fuel with greater temperature increase produces more energy NOTE: Fuels can be solids or liquids, for solid fuels we repeat the same steps but we hold the solid fuel below the test tube instead of in a spirit burner, and we use same mass and surface are of fuel instead of same volume 58  Investigation 11. Metal extraction and calculating percentages Questions ask you to extract a metal from its ore (e.g. Zn from ZnCO3, Cu from CuCO3, Sn from SnO) and calculate the percentage of metal in this ore. There are two methods to consider: a. Reduction - weigh a specific mass of the ore (initial mass) lid crucible - grind using morter and pestle - carry out thermal decomposition to obtain oxide - heat with carbon in a crucible heat - weigh the solid formed - percentage of metal= (mass of solid formed/ initial mass)* 100 b. Displacement - weigh a specific mass of the ore - grind using morter and pestle - add to nitric acid - add magnesium (more reactive metal) - filter to get the residue - wash and dry the residue - weigh the residue - percentage= (mass of residue/ initial mass)* 100 59  Practical Investigation 12. Preparing an indicator Questions as you to plan an experiment to prepare an indicator from a substance (e.g. fruit). Our goal is to simply extract the colour from the fruit. - Crush the fruit with pestle and mortar - add hot water and mix - then filter off, the filtrate is the solution of an indicator NOTE: We use the same steps for extracting the color from leaves TO SHOW THAT AN INDICATOR IS EFFECTIVE: - add the indicator to hydrochloric acid - add the indicator to sodium hydroxide - if it produces different colors with acid and alkali, then it can be used as an indicator (it is effective) 13. Chromatography The starting substance may be in an inappropriate form for chromatography, so we need to make it suitable before carrying out chromatography. To obtain a mixture of amino acid from protein and identify the amino acids - hydrolyze the protein by heating in concentrated hydrochloric acid - mixture of amino acids formed - use mixture of amino acids (this is our sample) - carry out chromatography (with steps) - use a locating agent (as amino acids produce colourless spots) - calculate the Rf value of each spot 60  Investigation MAIN CHROMATOGRAPHY STEPS: - draw baseline on chromatography paper by pencil - using glass rod add drop of sample on baseline - put paper chromatography in solvent (level of solvent must start below the sample) - leave it for a certain time - spots start to appear 14. Determining the solubility of a salt The solubility of a salt is the mass in grams of the salt that dissolves in 100 cm3 of water at a particular temperature; So we need to plan an experiment to determine this mass. - measure 100 cm3 of water using measuring cylinder - heat to 40oC - add salt until no more dissolves - stir - filter the mixture - evaporate filtrate to dryness - weigh the solid using balance OR - weigh a known mass of the salt (initial mass) - measure 100 cm3 of water using measuring cylinder - heat to 40oC - add the salt until no more dissolves - stir - filter to get the residue (undissolved salt) - weigh the residue (mass of undissolved salt) - mass of dissolved salt = initial mass - mass of undissolved salt 61  Practical Investigation 15. Separating two solids that react differently & calculating percentages Questions often mention that there is a mixture of two solids X & Y, X reacts with solution Z whilst Y does not, and want you to calculate the percentage of one of them. Our target here is to allow the reaction to take place then carry out our calculations. E.g. carbon and copper oxide are two black solids where only CuO reacts with H2SO4 To get the percentage of copper oxide & carbon - weight the mixture - add excess sulfuric acid - warm and stir - filter, wash and dry the residue (carbon) between filter paper - weigh the residue (carbon) - percentage of carbon = (mass of residue/mass of mixture) *100 - percentage of CuO = (mass of mixture-mass of residue/mass of mixture) *100 16. Thermal decomposition Questions might ask you to compare which substance produces more gas. Simply, we will carry out thermal decomposition and compare the results e.g. comparing CaCO3 and CuCO3 - set the apparatus as shown - place 20g of copper carbonate in the boiling tube - heat the metal carbonate in the boiling tube Copper carbonate gas syringe - reweigh - reheat till constant mass - measure the volume of gas produced - repeat the experiment with same mass of calcium carbonate - compare the volume of gas produced 62  Investigation 17. Maximum volume of oxygen that reacts with a metal Questions often want you to determine the maximum mass of oxygen that reacts with a *certain mass of metal granules* produicng metal oxide. Our target here is to allow the reaction to take place and measure the change in mass to get the maximum mass of oxygen. - use mortar and pestle to crush metal granules - add the powdered metal to crucible covered with lid - start heating, metal oxide is produced - lift the lid to allow entry of air - stop the reaction - reweigh the metal oxide formed -We crush the metal granules to increase - re-heat till constant mass surface are - maximum mass of oxygen = mass of metal -we cover the crucible with oxide – mass of metal a lid to prevent loss of metal oxide smoke 63