C: Practical Experiments :)

Mod 7 IQ5 Preparation of an Ester:

- Aim: To prepare an ester using reflux - Equipment § 10mL of Pentan-1-ol § 12mL of Glacial Acetic Acid § 1mL of conc. Sulfuric acid (CATALYST) § Boiling chips § 1-15mL of water § 15mL 1M Sodium Carbonate - Safety § Wear safety glasses § Concentrated sulfuric acid and acetic acid are corrosive. Handle with care. Avoid contact with eyes. If contact occurs, was with plenty of water. § Pentan-1-ol and pentyl-1-ethanoate are flammable and so, no naked flames should be present around them. 1. The refluxing apparatus with retort stand and condenser was set up and connected to a lab tap. 2. 10mL of pentan-1-ol, 12mL of glacial acetic acid (carboxylic acid) and 1mL of concentrated sulfuric acid (dehydrating agent gets rid of water to increase yield, is also a catalyst that is in excess to shift reaction right) were added to a 50mL round flask. 3. Approximately 10 boiling chips were added to the reaction flask and the flask was attached to the bottom of the condenser. 4. The tap was turned on to allow a uniform flow through the outside of the condenser (everything in condenser). 5. A heating mantel was placed underneath the flask switched on (collision theory). 6. After 30 minutes, the flask was left to cool for 5 minutes and then removed from the refluxing apparatus, and poured into a separating funnel containing 10-15mL of water. 7. A stopper was placed on top of the separating funnel, shaken, left to settle and drained from the bottom to remove the bottom aqueous layer (alcohol + water leaves only organic layer). 8. 15mL of 1M sodium carbonate was added to the solution left, stoppered, shaken through, left to settle and the bottom later was removed again from the solution (left over acid removed so ester is left). 9. The final substance was smelled carefully and observations recorded. - The original reagents react in equimolar quantities, so what is the purpose of adding one reagent in excess? To ensure that all of one reagent is used in the reaction.. Also ensures that a known amount of product is produced. - If the esterification reaction did not go through to completion, the final product could still be contaminated with reactant. How could you purify it further. By using fractional distillation, the substances within the mixture can be separated one by one to ensure greatest possible purity is achieved. - To isolate the ester (purify it), we should wash with water to remove soluble substances, then use a separating funnel to remove immiscible substances, then add sodium carbonate to remove any carboxylic acids.

Mod 5 IQ2 Cobalt (II) Chloride (Dehydrated/Hydrated):

- Cobalt chloride blue - Hydrated cobalt ions pink 1. Part 1 - Solution of hydrated cobalt (pink) when temperature increased, turned blue. 1. Water (approx. 1mL) was added using a pipette to 15mL of cobalt (II) chloride-alcohol solution (pink) until the solution just turns pink. 2. The solution was distributed equally into two test tubes and a stopper was placed over both tubes. One tube is the control. 3. The test tube not being used as a control was placed in 60oC hot water. Colour changes were observed. (turned from pink to blue) 4. The test tube that was placed in hot water was hen cooled under tap water and placed in an ice bath for 15 minutes. 5. The colour of the cooled test tube was compared to the control. (returned to pink) - Favours endothermic reaction (LCP - temp. change). 2. Part 2 - Filter paper (blue, dried) turned pink when water added. - Add 1 drop of water to the dried cobalt chloride filter paper and the colour change was observed. (turned from blue to pink) - The filter paper was then rested on a retort stand's clamp above a hot plate and dried again. Any changes were recorded. (restored blue colour). - Favoured endothermic (reverse) reaction (LCP - temp. change). - CoCl42-(aq) + 6H2O(l) ⇌ Co(H2O)62+(aq) + 4Cl-(aq) Blue Pink - CoCl2 + H2O ⇌ CoCl2 . 6H2O - Reaction from left to right is exothermic (gives out heat). - Add heat reverse is favoured. - Closed system stopper on end. - Le Chatelier's Principle is the reason for changes - if a system in chemical equilibrium is subjected to a disturbance it tends to change in a way that opposes this disturbance

Mod 5 IQ2: Iron III and Potassium Thiocyanate (Dynamic Equilibrium):

1. 50 mL of distilled water was added into a 100mL beaker. 2. A clean pipette for each was used to add in 2mL of 0.10mol/L iron (III) nitrate solution and 2.0mL of 0.10mol/L potassium thiocyanate solution. It was mixed thoroughly, and observations recorded. 3. 5mL of the mixture was added to five identical clean test tubes with 2 test tubes as controls for the two solutions without mixing (7 test tubes total). 4. Test tube A is the control of mixed solution. (maroon/blood red) 5. Test tube B had 10 drops of the iron(III) nitrate solution added (dark blood red). 6. Test tube C had 10 drops of the potassium thiocyanate solution added (dark blood red) 7. Test tube D had 5 drops of 2mol/L NaOH solution (light yellow/transparent) --> iron ions taken out of solution. 8. Test tube E had 1 gram of solid potassium nitrate added (orange with white sediment) --> product shifts left, excess ions move it to the right) - Fe3+ + SCN- ⇌ FeSCN2+ - (Yellow/orange) ⇌ (maroon/blood red) - Dynamic equilibrium - LCP - Relates to concentration.

Mod 7 IQ3 Modelling Saturated Hydrocarbons - Single Bonds:

1. A molecule of methane and several molecules of Cl2 was made. 2. Photos of the methane and chlorine were taken as 'reactant' photos. 3. Predictions were made. 4. Photos were taken of the 'products'. 5. Conclusion was achieved. - Limitations: oversimplification and no representation of bond interactions. - Substitution reaction. - Methane + Chlorine after many rounds can go to carbon tetrachloride and hydrogen (no hydrocarbon).

Mod 5 IQ4: Solubility:

1. Add 2 drops of each solution. 2. Find general rules. - Remember NAGSAG. - Only used balanced equations. N A G S A G i m r u c r t m o l e o r o u f t u a n p a a p t i t t e u 1 e e 1 s m s s 7 S PMS (lead, mercury, silver) & CASTROBEAR (calcium, strontium, barium) G2 PMS (lead, mercury, silver)

Mod 8 IQ1 Cation Test:

1. Add 5 drops of each solution to a spot test plate 2. Add sodium chloride and make qualitative observations 3. Add drops of each solution to a new spot test plate 4. Add ammonia solution and make qualitative observations 5. If no precipitate formed, add sulfate/hydroxide 6. Use a flowchart to determine presence based on observations. 7. Complete flame tests for final confirmation of presence Silver Ag+ No Distinct Flame Colour +Cl- -> White Precipitate +NH3 -> Brown Precipitate Lead (II) Pb2+ White/Grey +Cl- -> White Preciptate +NH3 -> White Precipitate Barium Ba2+ Apple Green +Cl- -> NP +NH3 -> NP +SO42- -> White Precipitate Calcium Ca2+ Brick Red +Cl- -> NP +NH3 -> NP +SO42- -> White Precipitate Copper (II) Cu2+ Blue-Green +Cl- -> NP +NH3 -> Blue Precipitate Iron (II) Fe2+ Golden Yellow +Cl- -> NP +NH3 -> Pale Green Precipitate Iron (III) Fe3+ Brown-Yellow +Cl- -> NP +NH3 -> Brown Precipitate Magnesium Mg2+ No Distinct Flame Colour +Cl- -> NP +NH3 -> NP +OH- -> White Precipitate Add: 1. HCl 2. NH3/NH3 3. - / Flame Test 4. - / NaOH - Ag (white, brown), Pb (white) / Cu (blue), Fe2+ (green), Fe3+ (brown) / [white) Ba (green flame), Ca (red flame) / Mg (white)

Mod 8 IQ1 Anion Test:

1. Add 5 drops of each solution to a spot test plate. 2. Add 3 drops of silver nitrate to each solution and make qualitative observations 3. If a precipitate forms, add excess nitric acid and make qualitative observations on whether the precipitate dissolved or remained. If no precipitate forms, add 5 drops of solution to new spot on spot test plate and add 3 drops of barium chloride and make qualitative observations on whether a precipitate formed. 4. Use a flowchart to determine presence based on observations. Sulfate SO42- Ag+ -> NP +BaCl2 -> forms white precipitate Chloride Cl- Ag+ -> white precipitate -> does not dissolve in HNO3 -> dissolves in dilute ammonia solution Bromide Br- Ag+ -> cream precipitate -> does not dissolve in HNO3 -> dissolves in concentrated ammonia solution Iodide I- Ag+ -> yellow precipitate -> does not dissolve in HNO3 Carbonate CO32- Ag+ -> white precipitate -> dissolves in HNO3 Phosphate PO43- Ag+ -> yellow precipitate -> dissolves in HNO3 Hydroxide OH- +Litmus -> blue + Cu(NO3)2 -> white precipitate Acetate CH3COO- +Litmus -> blue + Cu(NO3)2 -> NP Natural vinegar smell Add: 1. AgNO3 2. BaCl2 / HNO3 3. - / - / litmus - SO4 (none, white) / Cl (white), CO3 (white dissolves), Br (cream), I (yellow), PO4 (yellow dissolves) / OH- (red to blue) & CH3COO- (blue to red or vinegar smell) can be determined with litmus

Mod 7 IQ2 Bromine Water:

1. Bromine water was placed into three test tubes to a depth of 3cm. 2. Cyclohexane was added to two of the test tubes. Cyclohexene was added to the third test tub. Each was added to a depth of 1cm and test tubes were labelled. 3. A stopper was place on each tube. The tube was shakes and observations recorded. 4. One test tube containing cyclohexane was placed in a cupboard, and one test tube containing cyclohexane was placed in the sunlight. 5. Changes were noted, and appearances compared. - Cyclohexene was colourless. - Cyclohexane in UV was lighter yellow, not in UV was orange coloured both with dark brown/red upper layer. - Cyclohexene is addition reaction so it's more reactive. - UV adds energy to the molecules causing the reaction rate to speed up as outlined by collision theory.

Mod 6 IQ3 Titration:

1. Burette was rinsed with distilled water and then with HCl solution. Rinsing were discarded. 2. Burette was set up using a retort stand and burette clamp. 3. The stopcock was closed and the burette was filled with 50mL HCl. Initial volume was recorded. 4. 25mL Pipette was rinsed with NaHCO3 solution and rinsing discarded. 5. 25mL Pipette was filled with NaHCO3 solution and placed in clean 250mL conical flask. 3 drops of indicator was added (methyl orange). 6. Flask was placed under burette and HCl was run into flask while swirling until the colour changed. The end volume was recorded when this happened. 7. The volume of HCl used was recorded and steps 4-7 were repeated until 3 results were achieve within 0.1 of each other.

Mod 6 IQ3 Effect of Buffers:

1. Measure 10mL 0.2M acetic acid and pour in clean dry beaker. 2. Measure 40mL 0.2 mol sodium acetate and add to acetic acid (same) beaker. 3. Measure 20mL of this solution with measuring cylinder and place in a beaker. Repeat once. Discard remainder solution. 4. Add 5 drops universal indicator to both beakers. Record colour and pH. Label these beakers 'buffered solutions'. 5. In another beaker, add 50mL distilled water. 6. Add 10 drops of universal indicator to the water. 7. Add HCl drop by drop to distilled water until the colour is same as buffered solutions. 8. Measure 20mL of solution from step 7 with measuring cylinder into beaker (x2) and discard remainder. Label these unbuffered solutions and record colour. 9. Add 40 drops of 0.1 M HCl to one buffered solution, stirring continuously. Record colour and pH. 10. Repeat step 10 for one of the unbuffered solutions. 11. Repeat steps 9-10 with remaining solutions, adding 0.1M NaOH instead of HCl. - 2 buffered (certain pH) & 2 unbuffered (same starting pH as buffered by adding HCl). - Buffer helps find precise pH readings. - To improve, use a probe. - Buffer limits the change in pH making the shift more gradual and obvious in experiments. - Buffer is composed of a weak acid and conjugate base, so an equilibrium forms, to resist pH change. - H2O + CO2 ⇌ H2CO3 ⇌ HCO3- + H+

Mod 7 IQ4 Data Analysis of Boiling Point of Alcohols:

1. Molecular weight of alkanes and alcohols was calculated. 2. A chemical data book was used to determine the boiling points. 3. A graph was constructed to compare alkanes to alcohols. - Alcohols had higher boiling point. - Both linear trends - As chain length increases, the boiling point of alcohols and alkanes get closer together, but both increase steadily.

Mod 7 IQ4 Oxidation of Alcohols:

1. Prepare a hot water bath in a 250mL beaker. 2. Measure 2mL of each alcohol into a separate, labelled test tube. 3. In a separate test tube, mix 4mL potassium permanganate is purple with 1mL sulfuric acid. Note the colour of the solutions. 4. Add 4-5 drops of the acidified permanganate (oxidant) ion solution into each of the test tubes containing the alcohols. 5. Place the test tubes into the water bath for 1-2 minutes or until a change is seen. 6. Record any observations made, including colour changes. 7. Dispose of the contents of the test tubes in a waste container. Do not dispose of any materials down the sink. - Primary alcohol colour change from purple to clear when mixed with acidified permanganate (oxidises) - Secondary alcohol colour change from purple to clear when mixed with acidified permanganate (oxidises) - Tertiary alcohol colour change from purple to pale brown when mixed with acidified permanganate (does not oxidise) - Ethanol colour change from purple to clear when mixed with acidified permanganate - Primary and secondary are the same as there is available hydrogen bonded to the carbon bonded to the hydroxyl group which undergoes reaction, but the tertiary alcohol is unable to undergo the same reduction reaction due to the lack of the hydrogen.

Mod 6 IQ3 Digital Titration - Acid Content in Orange Juice:

1. Rinse 150mL beaker with sodium hydroxide solution (standardised NaOH about 0.1M). Empty, label and fill with 100mL NaOH solution. 2. Prepare burette (rinse it) then fill with NaOH solution. 3. Rinse the 250mL conical flask with water. 4. Measure the initial and final volume of NaOH. 5. Pipette 25mL of orange juice into conical flask and fill with 50mL distilled water. 6. Place pH meter into vessel, recording the pH. 7. Place conical flask under burette and begin titration while recording pH on software (continue until pH is consistently basic). 8. Repeat experiment 3 times and average results. - H3A + 3NaOH Na3A + 3H2O - Citric acid is triprotic - Steady flow in conical flask - Found rate of dispersion from graph (mL/sec) equivalence determined (no. moles used via concentration) - Titrant runs into the analyte (tube to beaker)

Mod 5 IQ1: Burning Steel Wool (Static Equilibrium) Prac:

1. Roll steel in ball & measure mass. 2. 9V battery. 3. Left react/cool. 4. Measure mass. - Fe2+ + O2 (combustion) - Ea too high to reverse - Non-reversible reaction (non-equilibrium) 2Fe(s) + O2(g) -------------> 2FeO(s) Heat Modelling Dynamic Equilibrium: 1. 2 measuring cylinders (A 100mL H2O, B 0mL) 2. Pipette (10mL) Water A-->B 3. Pipette (2mL) water B-->A - 2+3 = 1 transfer cycle 4. Volume water was measured in each after transfer cycle. 5. Repeated for 30 cycles. - Dynamic equilibrium - A --> 20 mL - B --> 80mL

Mod 7 IQ4 Enthalpy of Combustion of Alcohols:

1. Set up equipment as shown in above photo. 2. Put 100mL of water in beaker. 3. Measure initial mass of burner and initial water temperature and record in table below. 4. Put methanol burner below beaker and light. 5. Record change in temperature and final mass of burner. 6. Repeat for other alcohol burners. - Controlled variables - distance of flame, identical equipment needs to be managed. - Beaker will already be heated for following experiments so make sure it is cooled down before reuse. - Thermometer should not touch beaker. - Not stirring as this may speed up heating with atmosphere (as may introduce variable of stirring different amounts), but not stirring may mean it is not uniform temperature. - 11cm gap from base of beaker to retort stand, 1cm gap between beaker and burner. - Effect on equipment errors will be less significant with bigger temperature increase change. Lack of precision in equipment will lessen the impact with bigger temp change. - Fuel with the highest molecular mass will have higher molecular heat of combustion. - Enthalpy of Combustion Calculation: § Mass fuel = § ∆T = § q=mc∆T = X kJ § MM(Compound) = Y g/mol. § n=m/MM = mol. § ∆H = kJ/mol. = Z kJ/mol. = q/mol - The theoretical enthalpy of combustion of: 1. Methanol: -715.0 kJ/mol 2. Ethanol: −1366.8 kJ/mol 3. Propanol: -2021 kJ/mol 4. Butanol: - 2676kJ/mol. - Increases at steady rate (approx. 650 each time) - Ours is slightly lower as heat was lost to the surrounds.

Mod 8 IQ1 Flame Test:

1. Solutions of various salts are provided in atomiser bottles. 2. Place old newspaper on the bench. Place a lit Bunsen burner on the paper. Carefully spray the salt solution through a blue Bunsen burner flame and note the colour produced. 3. Copy the following table and record your observations. - Risk of heavy metal as it is toxic to humans such as iron and can cause damage to the body if ingested. Spray away from us and wash everything after. Water Orange/Yellow Copper Sulfate Cu2+ Green (Blue Green) Copper Chloride Cu2+ Green (Blue Green) Barium Chloride Ba2+ Light yellow (Apple Green) Barium Nitrate Ba2+ Light yellow (Apple Green) Calcium Chloride Ca2+ Orange (Orange Red) Calcium Nitrate Ca2+ Orange (Orange Red)

Mod 7 IQ3 Modelling Unsaturated Hydrocarbons - Double or Triple Bonds:

1. The addition of hydrogen to ethene was modelled. 2. Model the addition of chlorine, water and hydrogen bromide to ethene. - Remember naming rules - Benefits: visual representation of molecular sized process which is often invisible to the human eye. - Limitations: oversimplification and no representation of bond interactions.

Mod 6 IQ1 Testing Using Indicators:

1. You will be provided with 4 indicators; - Litmus - Phenolphthalein - Bromothymol blue - Methyl orange 2. Using a spot test plate observe and record the colour changes that occur when each of the above indicators is added to a solution of strong acid, weak acid, neutral solution, weak alkali and strong alkali.

Accuracy

Accurate · Manual vs digital tools · Scale of the measurement

Mod 6 IQ1 Measuring the Enthalpy of Neutralisation:

Aim: To determine the enthalpy of neutralisation and the effect of the state of reactants. Risks: NaOH is caustic, HCl is corrosive. Use spatula to transfer NaOH(s). Use safety glasses, personal protective clothing. Dispose of chemicals as directed to protect from splashing chemicals/contact with skin. Operate near water source (shower) in case of skin contact. Part A 1. Pour 100mL 1.0 mol/L HCl into a polystyrene cup. 2. Measure the temperature of this solution using the thermometer. 3. Accurately weight out approximately 4.0g of NaOH(s) and add this to the same polystyrene cup. 4. Use the thermometer to stir and record the highest or lowest temperature reached. Initial Mass NaOH 3.98g Volume HCl (1M) 100mL Initial Temp 26C Final Temp 51C Temp Change +25C - q = -mc∆T - = 100g x 4.18 x 25 - = 10450J - = 10.45 kJ - ∆Hneut = -104.5kJ (after being divided by 0.1 to account for the water) - If you want molar enthalpy, divide q by number of moles of H2O. Part B 1. Pour 50mL of 2.0mol/L NaOH in a polystyrene cup. 2. Measure the temperature of the NaOH using the thermometer. 3. Measure 50mL of 2.0mol/L HCl and record its initial temperature. 4. Average the initial temperatures of the NaOH and HCl (this is initial temperature of experiment). 5. Pour HCL into the polystyrene cup containing NaOH, stir and record highest or lowest temperature. Volume NaOH (2M) 50mL Volume HCl (2M) 50mL Avg. Initial Temp 26C Final Temp 40C Temp Change +14C - q = mc∆T - = 100g x 4.18 x 14 - = 5852 J - = 5.85 kJ - ∆Hneut = -58.5J

Mod 6 IQ1 Making a Natural Indicator:

Aim: To prepare an indicator solution from red cabbage and test the resulting indicator on a range of substances. Risk: Wear safety goggles (NaOH caustic, HCl corrosive). Avoid skin contact and wash immediately if touched. 1. Place shredded cabbage in 500mL beaker and just cover with distilled water (about 200mL). Slowly boil cabbage leaves until water turns dark reddish purple and leaves lose most of their colour. 2. Allow to cool and pour liquid into clean 250mL beaker. 3. Place 2mL of NaOH and HCl in separate test tubes. Add a few drops of red cabbage indicator until a definite colour is observed. Record colour. 4. Repeat step 3 with other substances and record results. Classify the substances as acidic (<6), basic (>8) or neutral (6-8). - Indicator turned pink/red for acids. - Indicator turned light green for bases.

Mod 6 IQ3 Making a Primary Standard:

General Risks: - Wear safety glasses as NaOH is corrosive and can cause eye damage. 1. Weigh out and measure accurately 1g NaHCO3 into a weighing boat. 2. Add 20mL distilled water and mix in a 20mL beakers, add to 250mL volumetric flask. Fill 2/3 of way with water, put stopper on and shake. Fill rest of way (pipette). Label and leave with stopper on top. - Primary stands (high purity, stability, high molar mass, known concentration, anhydrous [no water otherwise it is dilute])

Reliability

Reliable - similar results · Repeat at least 5 times · Range accepted error is +-0.2 · Average (remove error)

Mod 5 IQ3: Keq Iron (III) Thiocyanate Colourimetry:

Risks: - Chemicals splash in eyes/skin - safety glasses and wash hands before/after experiment. - FeSCN2+ harmful to aquatic life - don't pour down sink (waste bottle instead) Part 1 - 1. Label 6 volumetric flasks/test tubes. 2. Used 25mL bulb pipette to transfer 25 mL of 0.2mol/L Fe(NO3)3 to flask A. 3. Used a graduated 10mL pipette to transfer 1mL of 0.002mol/L KSCN to flask A. 4. HNO3 was added to make final volume of 100mL. 5. This was repeated 5 times with different KSCN and HNO3 values. 6. Cuvette was rinsed & colourimeter calibrated. 7. Cuvette filled ¾ with standard solution A and absorbance was measured & recorded in colourimeter. 8. Step 7 was repeated for all other solutions. (each solution was different colours - increased conc. = increased colour = increased absorbance). 9. Results graphed (conc. (x) and absorbance (y)). Part 2 - 1. Use equation of line to find concentration from absorbance. 2. Keq was found using concentration of solutions. - Fe3+ + SCN- ⇌ FeSCN2+ - (Yellow/orange) ⇌ (maroon/blood red)

Mod 6 IQ1 Demonstrating the Use of pH to Indicate the Difference Between the Strength of Acids and Bases:

Risks: - Wear safety glasses. - Carry one glassware with liquid for trip. - Wear protective clothing. Method: 1. Pour 25ml of 0.1M HCl, HNO3, acetic acid (CH3COOH), H3PO4, KOH, NaOH, ammonia (NH3) and NaF into separate test tubes and label as small beakers 1-7 respectively along with the name of solution it contains. 2. Submerge the pH meter into test tube 1 and record the pH shown on the pH meter. 3. Wash the pH probe/meter thoroughly with distilled water.* 4. Repeat steps 5-6 to measure the pH of a different substance in another test tube. *NOTE: You can also insert the pH probe in a buffer solution with a pH of 7 (i.e. neutral buffer) in a separate beaker to reset the pH reference point to 7 rather than rinsing the probe with distilled water. Note that distilled water has a pH of 7 which we used to rinse the pH probe. Results - Explanation: In the condition of equal acid molecules concentrations, compared to strong acids, weak acids will have a higher pH. This is because strong acids have greater concentration of hydrogen ions in solution due to complete dissociation/ ionisation. This would mean that strong acids have higher concentration of hydrogen ions in solution than weak acids provided the strong and weak acids have equal concentration to start with. In the condition of equal base molecule concentrations, compared to strong bases, weak bases will have a lower pH. This is because strong bases have more moles of hydroxide ions in solution if both the strong and weak bases due to 100% dissociation/ ionisation into hydroxide ions and its conjugate acid. This would mean that strong bases have higher concentration of hydroxide ions in solution than weak bases provided than the strong and weak bases have equal concentration. Acids Experiment Results: · HCl and HNO3 are strong acids. · CH3COOH and H3PO4 are weak acids. You will expect these to have higher pH (less acidic) than HCl and HNO3 since the concentration of the strong and weak acids are the same. Bases Experiment Results: · KOH, NaOH are strong bases. · NH3 is a weak base. You will expect these weak bases to have lower pH (less basic) than KOH and NaOH since the concentration of the strong and weak bases are the same. Most Acidic: · HCl & HNO3 - (strongest acids) · H3PO4 · WEAK CH3COOH Most Basic: · NaOH · KOH · WEAK NH3 / NH4OH

Validity

Valid · Variables · Control · Equipment · Repeat

How to determine reliability of a secondary source

· Comes from a reputable source (authority, and trustworthiness) · Look for it to be produced by more than one author · Needs to be current

VERBS:

· Discuss Identify issues, for and against. · Explain Cause & Effect · Assess Advantages, disadvantages, judgement · Evaluate describe issue, for and against, judgement based on criteria/question · Justify points supporting argument, conclusion (restate questions)

How to determine Accuracy of a secondary source:

· See the same information in more one place

How to determine Validity of secondary source:

· Whether or not the evidence provides the answer to the hypothesis

ACCURACY:

• Accuracy: Emphasis on exactness • The accuracy is the closeness to the true value. • In order to determine accuracy, the 'true value' or 'reference data' is needed

Mod 8 IQ1 Identifying Carbon-Carbon Double Bonds:

• Bromine Water Test • Hydration to form alcohol Test

LIMITATIONS OF MODELS:

• Can oversimplify processes • Can lead to a tainted understanding of a process eg. photosynthesis

Mod 8 IQ1 Identifying Carboxylic Acids

• Litmus test • Esterification Test

Mod 8 IQ1 Identify Hydroxyl Groups:

• Oxidation Test • Addition of Sodium produces hydrogen gas • Esterification Test

RELIABILITY:

• Reliability: emphasis on stability • Reliability is the repeatability or reducibility of the measurements. • Reliable measures are those that are similar in value over multiple experiments (conducted under the same conditions)

BENEFITS OF MODELS:

• Simplify processes to be easily understood • Allows processes to be visualised better • Can add colour to show clear differences • It can show atomic level reactions that cannot be seen by the naked eye

VALIDITY:

• Validity: Emphasis on meaning • Whether the experimental design/procedures measure what was intended to be measured • Accuracy, precision and reliability are all required for an experiment to be valid. • Experimental outcomes that are internally valid allow researchers to establish causation (a cause and effect relationship between the independent and dependent variables) • Results that are externally valid allow researchers to predict future outcomes and extrapolate current findings to new conditions