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Presentation On Chemistry AS/A Level

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Published in: Chemistry
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Chemistry AS/A level

Areesha A / Dubai

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Qualification: IGCSE-AS Level-A level

Teaches: Biology, Chemistry, English, Physics, Science, Phonics, Maths, English Language

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  1. 1.4: Empirical and molecular formulas, using significant figures LESSON OBJECTIVE: Understand and calculate empirical and molecular formulas. Understand how to report calculations to the correct amount of significant figures. Learning Outcomes: (taken from the Cambridge International AS and A Level Chemistry (9701) 2019-2021 curriculum) 1.4 The calculation of empirical and molecular formulae a) define and use the terms empirical and molecular formula b) calculate empirical and molecular formulae, using combustion data or composition by mass C) Understand and use the terms anhydrous, hydrated and water of crystallisation. https ://www.youtube.com/watch?v—Ou FqtxZJRvM https://edu.rsc.orq/resources/the-chanqe- in-mass-when-maqnesium- burns/ 71 8.article Resource Plus Carry out the Finding the empirical formula by displacement experiment referring to the Teaching Pack for lesson plans and resources.
  2. Empirical and Molecular formula Empirical formula: the formula that tells us the simplest ratio of the different atoms present in a molecule Molecular formula: the formula that tells us the actual numbers of each type of atom in a molecule MOLECULAR p 4010 C02 N204 C5H120 EMPIRICAL p 205 C02 N02 Cl-12 C5H120
  3. What is the empirical formula for the following? Hydrazine, N2H4 Benzene, C6H6 - NH2 Octane C H 8 18 - C4H9 Ammonia, NH3 - NH3
  4. Percentage mass by composition atomic mass x number of moles of particular element in a compound % by mass = molar mass of compound Calculate the percentage by mass of iron in iron (Ill) oxide. M(Fe) = 55.85, M(O) = 16.0 % mass (Fe) = 69.9% (3 s.f) x 100%
  5. Significant figures and rounding The amount of significant figures needs to be consistent with the data given. The significant figures imply the level of precision the data was measured to. Your answer can only be as accurate as the least precise measurement of data that you have. Need to round to the correct number of significant figures but ONLY ROUND IN THE VERY LAST STEP OF YOUR CALCULATIONS Rounding to early can result in a rounding error Worked example: How many moles are there in 4.0g of potassium oxide?
  6. Determining the empirical formula Method 1: by combustion. Example - burning magnesium. Step 1: Burn a known mass of magnesium e.g. 0.486g Step 2: Record the mass of the product after combustion (i.e. magnesium oxide) e.g. 0.806g Step 3: Calculate the number of moles of the magnesium and the oxygen use n = m/ How do we know the mass of O? Step 4: Divide each by the lowest figure to get a simple ratio. Write the empirical formula. Mg o
  7. Determining the empirical formula Method 2: by percentage mass. Example, a hydrocarbon compound contains 80% Carbon and 20% Hydrogen. c Step 1: Note the % by mass. Step 2: Divide by the Ar values. Step 3: Divide by the lowest figure to determine the ratio. Write the empirical formula. H
  8. Calculating the molecular formula Determine the molecular formula from... The relative molecular mass (Mr) The empirical formula mass E.g. A compound has the empirical formula CH2Br and 187.8, deduce the molecular formula. Step 1: Find the empirical formula mass (add up each of the atoms molar mass). Step 2: Divide the relative molecular mass by the empirical formula mass. Step 3: Multiply the number of each atom by the answer in step 2.
  9. PAST PAPER (MCQ & STRUCTURED) Which volume of hydrogen, measured under room conditions, is produced when 0.160g of methanol reacts with an excess of sodium? A 60cm3 B 120 cm3 c 240 cm3 The respmses A to D should be selected on the basis of D 480 cm3 Compound O contains 40% carbon by mass. What could Q be? 1.2and3 1 and 2 only are 2 and 3 only are correct 1 only correct 1 2 3 glucose, C6H,206 starch, (CcH100s)n sucrose, C12H220" NO Other combinatbn Of statements is used as a correct response. In 1814. Sir Humphrey Davy and Michael Faraday collected samples Of a flammable gas, A, from the ground near Florence in Italy. They analysed A which they found to be a hydrocarbon. Further experiments were then carried out to determine the molecular formula Of A. (a) What is meant by the term molecular formula? . [21 Davy and Faraday deduced the formula Of A by exploding it with an excess Of oxygen and analysing the products of combustion. (b) Complete and balance the following equation for the complete combustion of a hydrocarbon with the formula C XH CxHy (x
  10. PAST PAPER (STRUCTURED) (c) When 10 cm3 Of A was mixed at room temperature with 50 cm3 of oxygen (an excess) and exploded, 40 cm3 of gas remained after cooling the apparatus to room temperature and pressure. When this 40 cm3 of gas was shaken with an excess of aqueous potassium hydroxide, KOH, 30 cm3 of gas still remained. What is the identity of the 30 cm3 of gas that remained at the end of the experiment? The combustion Of A produced a gas that reacted with the KOH(aq)_ What is the identity Of this gas? What volume Of the gas you have identified in (ii) was produced by the combustion (iv) of A? 3 What volume of oxygen was used up in the combustion of A? cm 3 141 (d) use your equation in (b) and your results from and to calculate the molecular formula of A Show all Of your working. [Total: 111
  11. PAST PAPER (MCQ & STRUCTURED) ANSWERS (a) the actual number of atorns of each element present (1) in one molecule Of a cornpound (I) (b) CxHy+ x + Y) 02 xC02 + xcoz (1) oxygen/02 ( I ) carmn dioxide/C02 ( 1 ) 10 crn3 (1) 20 (1 ) x + E 02 xC02 (d) CxHy _ H20 y H20 10 20 c: r•n3 I mol Of gives I mol Of C02 whencex= 1 (1) 1 mol of C.Hy reacts with 2 rnol of 02 and molecular forrnula is CH. (1) [2] [Total: 111
  12. WORK TOWARDS THE TOP Work through each task, one column at a time. Start at the bottom and work your way up. Calculate the percentage mass by composition of copper in copper (Il) oxide (CLIO). Give examples of the empirical and molecular formulas of certain molecules. Copy the definition for: Empirical formula Molecular formula START AT THE BOTTOM Research the limitations for determining the empirical formula via combustion, suggest ways you could overcome these in the lab. Explain how to calculate the molecular formula when you are given an empirical formula and a relative molecular mass. Summarise the two methods for determining the empirical formula of a compound. Investigate the e rounding error and suggest ways to avoid this. Create a document you could give to someone that explains how to determine the correct amount of significant figures. Summarise the importance of significant figures when undertaking calculations.
  13. LESSON OBJECTIVE: Understand and use the terms anhydrous, hydrated and water of crystallisation. Heating in a crucible This method could be used for measuring mass loss in various thermal decomposition reactions and also for mass gain when reacting magnesium in oxygen. The water of crystallisation in calcium sulphate crystals can be removed as water vapour by heating as shown in the following equation. CaS04.xH20(s) — CaS04(s) + xH20(g) Method. •Weigh an empty clean dry crucible and lid . •Add 2g of hydrated calcium sulphate to the crucible and weigh again •Heat strongly with a Bunsen for a couple of minutes •Allow to cool •Weigh the crucible and contents again •Heat crucible again and reweigh until you reach a constant mass ( do this to ensure reaction is complete). Large amounts of hydrated calcium sulphate, such as 50g, should not be used in this experiment as the decomposition is like to be incomplete. The crucible needs to be dry otherwise a wet crucible would give an inaccurate result. It would cause mass loss to be too large as water would be lost when heating. The lid improves the accuracy of the experiment as it prevents loss of solid from the crucible but should be loose fitting to allow gas to escape. Small amounts the solid , such as 0.100 g, should not be used in this experiment as errors in weighing are too high.
  14. Exarnple 3.51 g Of hydrated zinc sulphate w•ere heated and 1 -97 g Of anhydrous zinc sulphate were obtained. Use these data to calculate the value Of the integer x in ZnS04-xH20 Calculate the rnass Of H20 = 3.51 — 1.97 = 1.54g Calculate moles of Znsoa = 1.97 161.5 = 0.0122 Calculate moles of 1--120 = 1.54 18 —0.085 0.0122 Calculate ratio Of mole of ZnS04 to H20 Hazards and Risks - QUI.22 0.0122 A hazard is a substance or procedure that can has the potential to do harm. Typtcal hazards are toxic/flammable Iharmfub' irritant Icorrosive 'oxidizing/ carcinogenic In the laboratory we try to minimise the risk Irritant - dilute acid and alkalis- wear googles Corrosive- stronger acids and alkalis wear goggles Flammable — keep away from naked flames RISK: This is the probability or chance that harm will result from the use of a hazardous substance or a procedure Hazardous substances in low concentrations or amounts will not pose the same risks as the pure substance. Toxic — wear gloves- avoid skin contact- wash hands after use Oxidising- Keep away from flammable I easily oxidised materials
  15. 1.5: Thermal decom ositiono h drated cr stals LESSON OBJECTIVE: Understand and use the terms anhydrous, hydrated and water of crystallisation. Practical 1.1: Findin the formula of h drated CODDer sulfate crystals (fäkéiiTFöiiiThäCöfiibridge In}ernationa1AS and A Le el Chemistry (9701) 2019-2021 cur iculum) https://edu.rsc.org/resources/finding-the-formula-of-hydrated-copper-ii- sulfate/436.article Safety • Eye protection should be worn. • Copper(ll) sulfate is harmful if swallowed and is irritating to the eyes and skin. For disposal, dissolve 64 g of the anhydrous form in 1 litre of water before pouring down the foul-water drain; keep this disposal procedure to a minimum. Notes on the procedure • Give each group a different mass of copper(ll) sulfate crystals to heat, ranging from 0.25 g to 2.00 g. The results can then be pooled. • Teachers should not allow students to heat the copper sulfate so hard that the white fumes of corrosive sulfur dioxide and toxic sulfur dioxide are produced. A sign that this is happening is that the anhydrous copper(ll) sulfate discolours into red and black powders.
  16. 1.5: Thermal decomposition of hydrated crystals Practical 1.1: Finding the formula of hydrated copper sulfate crystals (taken from the Cambridge International AS and A Level Chemistry (9701) 2019-2021 curriculum) Ideal' results (from a 2-decimal-place balance) are given in the following table. Group Mass Of CuS04 crystals / g Mass Of anhydrous CuS04 / g Mass Of water / g 0.25 0.16 0.09 2 0.50 0.32 0.18 3 1.00 0.64 0.36 4 1 .25 080 0.45 5 1.50 0.96 0.54 6 1.75 1.12 0.63 7 2.00 1.2B 0.72 An overhead transparency showing the 'ideal' line for the experimental results might prove useful. There are two main sources of error in the experiment: •If students heat the copper(ll) sulfate too strongly and there is decomposition; in this case their point will lie below the 'ideal' line. •If students do not heat the copper(ll) sulfate enough, in which case their point will lie above the 'ideal' line.
  17. LESSON OBJECTIVE: Understand and use the terms anhydrous, hydrated and water of crystallisation. (taken from the Cambridge International AS and A Level Chemistry (9701) 2019-2021 curriculum) In this practical you will: • make a solution of washing soda • titrate the washing soda solution against standard hydrochloric acid • calculate the formula of washing soda crystals (Na2C03.xH20), using the titration results. Safety • Wear eye protection at all times. • Washing soda powder is an irritant if inhaled and methyl orange is toxic. Making a solution of washing soda • Apparatus (per student washing sodåOYSta s spatula weighing pot 250 cm3 beaker 250 cm3 volumetric flask, with stepper g ass stirring ro glass filter funnel distilled water access to accurate balance e e rotection Then add distiu a dm r the https://wmv.youtube.com/watch?v=lbkreR5VDmo
  18. • • Procedure 1 Accurately weigh out 1.29 g of the washing soda crystals (IRRITANT). 2 Carefully transfer the crystals into the 250 cm3 beaker, ensuring that all the powder is transferred by repeatedly washing the weighing pot and transferring the washings into the beaker. 3 Half-fill the beaker with distilled water. Stir until all the crystals are dissolved. 4 Place the funnel in the neck of the volumetric flask. 5 Transfer the solution to the volumetric flask. Wash the beaker with a few cm3 of distilled water and transfer these washings into the volumetric flask. 6 Repeat this a few times and then wash the funnel with a small volume of water. 7 Finally make up to the mark on the volumetric flask with distilled water. 8 Stopper the flask and then invert nine times so that a homogeneous solution is obtained.
  19. Titrating the washing soda solution against standard hydrochloric acid Apparatus (per student) 250 cm3 of standard hydrochloric acid, O. 100 mol dm-3 washing soda solution burette and burette stand 25 cm3 ipette and filler Calculation The equation for the reaction is: small funnel conical flask white tile screened methyl orange indicator e e rotection 1 2 3 4 5 Na2C03(aq) + 2HCl(aq) -+ 2NaCl(aq) + H20(l) + C02(g) (Ar values: Na 23.0, C - 12.0, H - 1.0, O - 16.0) Procedure Fill the burette with the hydrochloric acid. Pipette 25.0 cm3 of the washing soda solution into the conical flask. Titrate the hydrochloric acid against the washing soda solution, using screened methyl orange (TOXIC) as indicator. Construct a results table. Do one rough titration and then accurate titrations until concordant results are obtained. 1 Calculate the average titre (volume of hydrochloric acid) from your results. 2 Calculate the number of moles of hydrochloric acid present in this volume of acid. 3 Using the equation, calculate the number of moles of sodium carbonate present in 25 cm3 of solution. 4 5 Calculate the number of moles (n) of sodium carbonate (washing soda) crystals present in the 250 cm3 of solution in your volumetric flask Using the formula Mr where m is the mass of washing soda you used) find the relative molecular mass ogodium carbonate and hence the value of x in the formula
  20. MEASUREMENTS & ERRORS DEFINITIONS: •Accuracy is the closeness of agreement between a measured value and the true value. •Precision is the closeness of agreement between independent measurements of a quantity under the same conditions. •Uncertainty is the component of a reported value that characterizes the range of values within which the true value is asserted to lie. •Error is the difference between a measurement and the true value of the object being measured. Error does not include mistakes. Calculation of percentage error Percentage error = (maximum error * quantity measured ) x 100% Examples: Consider weighing lg of solid. If you use a two decimal place balance, the mass recorded will be to the nearest 0.01 g. We should express this measurement as (1.00 ± 0.01) g, where 0.01 is the absolute error. In this example, the % error will be: 0.01 g x 100 =
  21. MEASUREMENTS & ERRORS IN TITRATION rror •n ure es A burette is graduated in divisions every 0.1 cm3. Using the half-division rule, the estimation is 0.05 cm3. Burette is recorded to two decimal places with the last figure either co' or '5'. The maximum error in each measurement = 0.05 cm3. The overall maximum error in any volume measured always comes from two measurements, so the overall maximum error = 2 x 0.05 cm3 = 0.1 cm3. In a titration, a burette will typically deliver around 25 cm3 so the percentage error is small. In your practical exam, if they ask you to calculate the percent error in the titration, you should use the average of the volume you calculated in the titration as the volume measured. Example, you perform the titration, and you calculate that the average volume is 24.50 cm3 Percentage error = (2 x 0.05 cm3 +24.50 cm3) x 100% = 0.41% (no units) The percentage error becomes more significant when burette is used to deliver small volume For delivery of 5.50 cm3, Percentage error = (2 x 0.05 cm3 + 5.50 cm3)x 100% = 1.82% (no units) Error in 25 cm3 Pipettes The error for a 25 cm3 pipette, is written on the apparatus. Most of the time it will be around 0.02mL (cm3), so the error for the pipette will be absolute error: 0.02 cm3 Percentage error = ( 0.02 cm3 + 25.00 cm3)x 100% = 0.08% Since the error when measuring with the burette is much bigger, it will be the one affecting the results the most.

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