volatile liquid experiment leaving cert

Experiments​​

28 Mandatory Chemistry Investigations  

To carry out flame tests with different salts.

Test for anions in aqueous solutions: chloride, sulfate, sulfite, carbonate, hydrogencarbonate, nitrate, phosphate.

To measure the relative molecular mass of a   volatile liquid.

To prepare a standard solution of sodium carbonate.

To use the standard solution of sodium carbonate to standardise a given hydrochloric acid solution .

  5A  To prepare a 0.1M solution of sodium hydroxide and       standardise it with a standard hydrochloric acid solution to        prepare sodium chloride.

  6. To determine the concentration of ethanoic acid in vinegar.

  7. To determine the amount of water of crystallisation in                hydrated sodium carbonate (washing soda).

  8. To study some oxidation-reduction reactions with a)                  Halogens as oxidising agents and b) Displacement                    reactions of metals.

  9. To prepare a standard solution of ammonium iron(II)       sulfate and to use this solution to standardise a solution   of        potassium permanganate by titration.

 10. To determine the amount of iron in an iron tablet .

 11. To prepare a solution of sodium thiosulfate and to       standardise it by titration against a solution of iodine .

 12. To determine the percentage (w/v) of sodium       hypochlorite in household bleach.

 13 . To monitor the rate of production of oxygen from        hydrogen peroxide using manganese dioxide (catalyst).

14A To study the effect of concentration on the reaction rate       using sodium thiosulfate and hydrochloric acid.

14B To study the effect of temperature on the reaction rate       using sodium thiosulfate and hydrochloric acid.

 15. To illustrate Le Chatelier's principle using the reaction        between iron(III)chloride and potassium thiocyanate.

 16. To determine the total water hardness in a water sample           using EDTA.

 17. To determine (i) the total suspended solids in ppm of a             sample of water by filtration, and (ii) the total dissolved           solids in ppm of a sample of water by evaporation, and            (iii) the pH of a sample of water.

 18. To measure the amount of dissolved oxygen in a       sample of water by means of a redox titration.

 19. To estimate the concentration of free chlorine in a                     swimming pool using (i) a comparator and (ii) a                       colorimeter.

 20. To prepare ethene and examine its properties.

 21. To prepare ethyne and examine its properties.

 22. To determine the heat of reaction of hydrochloric acid with        sodium hydroxide.

23A To extract clove oil from cloves by steam distillation.

23B To isolate clove oil (eugenol) from an emulsion of clove       oil and water by liquid liquid extraction using cyclohexane.

 24. To prepare a sample of soap .

25A To study the reaction of ethanal with a acidified potassium         permanganate.

25B To study the reaction of ethanal with Fehling's reagent.

25C To study the reaction of ethanal with ammoniacal silver             nitrate (silver mirror test).

25D To study the reaction of ethanoic acid with (i) sodium               carbonate, (ii) magnesium and (iii) ethanol.

 26.   To oxidise phenylmethanol to benzoic acid with potassium         permanganate under alkaline conditions.

 27. To (i) recrystallise a sample of benzoic acid and (ii)                   measure it's melting point.

 28. To separate the components of ink using paper                         chromatography or thin-layer chromatography or column           chromatography.

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Orbitals overlapping head on is Sigma every time.

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Having problems seeing orbitals in 3 dimensions? Check out 2.4 Atomic Orbitals & Electronic Configurations.

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Over 80 animated videos that take you from first principles right through to everything you need to know for the exam. More than 400 handcrafted quizzes, modelled on past paper questions, that will test your understanding every step of the way. ‍ Over 80 sets of illustrated revision notes that summarise all of the key details you need to know. Our Past Paper Portal categorises all past paper questions by topic making it easy for you to find the questions you need so you can spend time doing them rather than searching for them.

Revision Notes

Countless past paper qs, how we’ve laid it all out.

1.1 The History of the Atom

1.2 Arrangement of electrons in the Atom

1.3 Radioactivity

ME 1.1 Flame Tests

2.1 History of the Periodic Table

2.2 Using the Periodic Table

2.3 Isotopes, Relative Atomic Mass & The Mass Spectrometer

2.4 Atomic Orbitals & Electronic Configurations

2.5 Trends in the Periodic Table

3.1 Ionic Bonding

3.2 Covalent Bonding

3.3 Shapes of Molecules

3.4 Electronegativity and Molecular Polarity

3.5 Intermolecular Forces

ME 3.1 Tests for Anions

4.1 Balancing Chemical Equations

4.2 The Mole

5.1 The Gas Laws

5.2 Kinetic Theory of Gases & The Ideal Gas Equation

ME 5.1 To Find the Relative Molecular Mass of a Volatile Liquid

6.1 Mass-Moles-Particles

6.2 Percentage Composition by Mass & Empirical Formulas

6.3 Mass-Mass Equations

6.4 Gases in Equations

6.5 Solutions

6.6 Solutions in Equations

6.7 Further Stoichiometry

7.1 Types of Acids and Bases

7.2 Neutralisation Reactions & Volumetric Analysis

ME 7.1 To Prepare a Standard Solution of Sodium Carbonate

ME 7.2 Standarising a HCl Solution

ME 7.3 To Determine the Concentration of Ethanoic Acid in Vinegar

ME 7.4 To Determine the Water of Crystallisation in Washing Soda

8.1 Oxidation and Reduction Reactions

8.2 Oxidation Numbers

8.3 Electrochemistry

ME 8.1 REDOX Reactions of a) Metals and b) Halogens

ME 8.2 To Standardise a Solution of Potassium Permanganate

ME 8.3 To Determine the Amount of Iron in an Iron Tablet

ME 8.4 To Standardise a solution of Sodium Thiosulphate against a solution of Iodine

ME 8.5 To Determine the Percentage of Sodium Hypochlorite in Bleach

9.1 What is Rate of Reaction?

9.2 Measuring Rates of Reactions

9.3 Average Vs Instantaneous Rates of Reaction

9.4 Collision Theory

9.5 Factors that affect Rate of Reaction

9.6 Catalysts

ME 9.1 Decomposition of Hydrogen Perioxide

ME 9.2 Sodium Thiosulphate + HCl

10.1 What is Chemical Equilibrium

10.2 The Equilibrium Law

10.3 Problems using Kc

10.4 Le Chatelier’s Principle

ME 10.1 Le Chatlelier’s Principle

11.1 Water and pH

11.2 pH Calculations

11.3 Indicators

12.1 Hardness in Water

12.2 Water treatment

12.3 Pollution

ME 12.1 To determine the total hardness in a water sample using EDTA

ME 12.2 To determine the total suspended and dissolved solids in a water sample

ME 12.3 To measure the amount of dissolved oxygen in a sample of water by means of a REDOX titration

ME 12.4 To estimate the concentration of free chlorine in swimming-pool water

13.1 Hydrocarbons

13.2 Refining of Oil

13.3 Thermochemistry

ME 13.1 To Prepare Ethene and Examine its Properties Summary Sheet.html

ME 13.2 To prepare Ethyne and examine its properties

ME 13.3 To determine the Heat of Reaction of Hydrochloric acid with Sodium Hydroxide

14.1 Tetrahedral Organic Compounds

14.2 Planar Organic Compounds

14.3 Substitution Reactions

14.4 Addition & Elimination Reactions

14.5 REDOX Reactions and Organic Acids

14.6 Organic Synthesis

14.7 Chromatography and Instrumentation

ME 14.1 Extraction of Clove Oil

ME 14.2 To Prepare a Sample of Soap

ME 14.3 Ethanal and Ethanoic Acid

ME 14.4 Oxidation of Phenylmethanol

ME 14.5 Benzoic Acid

ME 14.6 Chromatography

Crunch Chemistry

The ultimate A level chemistry resource

Finding the mass of a volatile liquid

We can use the ideal gas equation to find the molar mass (relative molecular mass) of a volatile liquid. In this video I briefly run through the experiment, the measurements we need and a sample calculation – foolproof method guaranteed!

So, you are now ready to tackle the exam style questions (relevant for all A level exam boards) 😀.

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Chemistry Experiment predictions

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Gas Laws, moles & Gas Properties

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General Chemistry Lab News

Middlebury college.

Identification of an unknown volatile liquid

Learning goals.

• Use bomb calorimetry to measure heat of reaction;
• Use a thermometer and analyze the limitations of thermometry;
• Observe changes in phases of matter and consider the challenges associated with volatile substances;
• Apply the ideal gas law and analyze the limitations of using the ideal gas law in an experiment;
• Synthesize a simple procedure and a data table to collect experimental data;
• Choose a follow-up experiment based on initial results;
• Write a conclusion based on a specific prompt and tabulate results to support the conclusion.

Ideas for portions of this lab were drawn from the University of Pennsylvania Laboratory Program for Chem 53, accessed at this webpage: http://www.sas.upenn.edu/~kennethp/chemlab1.pdf, and the University of Calgary Laboratory Program for Organic Chemistry, accessed at this webpage: http://www.chem.ucalgary.ca/courses/351/laboratory/boilingpoint.pdf.

Introduction

https://qph.ec.quoracdn.net/main-qimg-4161e9231033aa335aca697a3ef0c4e2

In many ways, chemistry can be a puzzle: since you cannot see atoms directly, you must instead collect as many pieces of data as you can and put them together to get the best possible answer. In this lab, you are tasked with identifying an unknown volatile liquid.

First, you will measure the molar mass of the liquid using the ideal gas law, but is the molar mass enough information to identify an unknown? How much uncertainty is associated with the molar mass measurement, anyway? If you were publishing a paper and wanted to convince the scientific community that you have identified a liquid, measuring the molar mass would surely not be enough information!

As additional pieces of evidence, you will measure the density in the liquid state, the boiling point, and the heat of combustion. These four pieces of evidence should pinpoint the exact identity of the liquid with enough certainty to convince even the most skeptical audience, who is hopefully yourself.

When you think you know the identity of the unknown, you can run a confirmatory test.

What is a volatile liquid?

A volatile liquid is a liquid that vaporizes readily under normal conditions. Many volatile liquids have a noticeable smell because they vaporize into your nose! Throughout this lab, you will be working with a volatile liquid. When it is warm, it is mostly a gas, and when it is cool it is mostly a liquid. Note that you will be measuring the density of the substance in both the liquid and the gaseous state, but you will find that the tendency of the volatile liquid to vaporize and condense introduces uncertainty in both of these measurements.

Measuring molar mass from density of the gaseous state

The ideal gas law provides a way to relate the number of moles of a gas to the volume of the gas: n/V = P/RT

If you are also able to relate the number of grams of the gas to the volume of the gas (g/L, in other words, density), then these two formulas together allow you to then relate the number of grams of the gas and the moles of the gas, which is the molar mass: molar mass = (density)RT / P.

Your textbook shows a more detailed derivation in Section 8.3.

In this lab, you will measure the density of a volatile liquid in its gaseous state, the pressure, and the temperature, and this will allow you to calculate the molar mass. To measure the density, you will need to measure the mass and volume of a gas. The procedure will guide you through forming the gas and measuring the mass, but you must design your own method of measuring the volume of the gas. Be creative! Try to get as many significant figures as possible.

Heat of combustion

This is a topic you have not yet covered, but you will soon! Here is a simple explanation of what you need to know for this lab:

The heat of combustion is the amount of energy released when a given amount of a substance is burned in the presence of excess oxygen. Normally this value is reported in terms of energy per mole of the substance, but because we do not yet know the molar mass of the liquid, we will measure it in terms of energy per gram of the substance and compare it to the list of possible sample identities at the bottom of this page, also in those units, in order to identify the liquid.

Density of the liquid state

You should be very comfortable measuring the density of a liquid. You just need the mass and the volume: density = mass/volume. You will be comparing the density of your liquid to the list of possible sample identities at the bottom of this page to help you identify the liquid.

Refer back to Lab 1 for a refresher on density measurements.

Boiling point

Boiling point is the temperature at which a substance changes from a liquid state to a gaseous state. Boiling point is actually a function of pressure, but because most people measure it under atmospheric pressure, it is typically reported at 1.000 atm (1013 mbar). If you notice that the atmospheric pressure on the day of your measurement is greater than or less than 1013 mbar, know that this will contribute to a discrepancy between your results and the literature boiling point values.

Note: do this lab with a partner. There are five parts to this lab, which will be done over two weeks. I suggest doing part 1 on the first week and parts 2-5 on the second week, although you may chose to do parts 2 and 3 whenever it is convenient. Be sure to use the same thermometer throughout the experiment.

1. Measuring molar mass from density

Thermometer calibration.

At some point, you probably had this thought but were afraid to say it out loud: “are these thermometers right? I always thought water boiled at 100 °C, but it looks like this water is boiling at 105 °C. I don’t know, maybe I’m reading it wrong?” If you have had this thought, then pat yourself on the back for being observant. The alcohol thermometers are not perfect, and to get a reasonably accurate temperature measurement with them, you will need to calibrate them. Since we are making most of our temperature measurements in this experiment near 100 °C, we will calibrate at 100 °C.

Bring a beaker of tap water to a rolling boil over a bunsen burner. It doesn’t matter what size beaker and exactly how much water, but a 600 mL beaker about 3/4 filled with tap water would be a great choice. Once boiling, place your thermometer into the water. Hold it or clamp it so it doesn’t touch the bottom or sides. When the temperature is stable, record the “thermometer calibration temperature” to the nearest 0.2 °C in Table 1.

You can now calculate the offset for your individual thermometer.

100.0°C – “thermometer calibration temperature”(°C) = thermometer offset (°C)

This number may be positive or negative. Use a piece of label tape to label your thermometer with its thermometer offset. From now on, factor that offset into your temperature measurements before recording them in your data table. Keep that calculator handy!

As an example, let’s say your thermometer calibration temperature is 105.0 °C.

100.0°C – 105.0 °C = -5.0 °C

Now you know to subtract 5.0 °C from every temperature you measure throughout the experiment. If your thermometer reads 92.2 °C, then record 92.2 °C – 5.0 °C = 87.2 °C.

Record the barometric pressure of the atmosphere in Table 1. If you have reason to believe the pressure is changing throughout the experiment (maybe there is a hurricane outside?), check it before beginning every trial, but normally checking at the beginning of the experiment is sufficient.

Mass and temperature

Use the same 125-mL Erlenmeyer flask throughout this section. Dry the Erlenmeyer flask thoroughly. Measure the mass of the 125-mL Erlenmeyer flask with an Aluminum foil cover secured with a rubber band. Trim the Aluminum foil cover to right below the rubber band. Record the mass of the covered flask in Table 1. Make a tiny hole in the foil cover with a thumbtack.

Bring the water to a boil over a blue bunsen flame.

While the water begins to heat, remove the foil cover and rubber band, place about 2 mL of the volatile liquid from the small bottle into the Erlenmeyer flask, then replace the foil cover and rubber band.

Turn off the flame. Submerge the covered Erlenmeyer flask up to its neck in the hot water, and use a clamp to attach it to the ring stand. Add more water if you need it to cover the flask up to the neck, but try not to get the aluminum foil wet.

Also, submerge a thermometer in the water, and use a clamp to attach it to the ring stand.

Monitor the temperature of the water to ensure it stays around 90 °C (don’t forget to factor in your calibration offset). Turn on the flame if the temperature drops below 80 °C. Maintaining this temperature, watch the vapor escape through the pinhole. Watch for the vapor to stop escaping through the pinhole and for visible condensation in the flask to disappear (tendrils dripping down the sides of the inside of the flask). This should take no more than 10 minutes. If you continue seeing condensation after the first 10 minutes, turn up the heat because your thermometer may be reading low. You can verify that the vapor has stopped escaping by holding a small watch glass over the pinhole and seeing that no condensation develops right above the pinhole. As soon as the vapor stops, record the temperature of the hot water bath to the nearest 0.2 °C in Table 1. Be sure to factor in the calibration offset!

Remove the Erlenmeyer flask from the hot water bath, but keep the cover on. Be sure all of the liquid has vaporized. If there is visible liquid, put it back into the hot water bath for a few minutes. Dry the outside of the flask completely. Then, measure the mass of the covered flask with the vapor inside (it may condense back to a liquid). If the mass is not stable, let it cool longer. Record the mass in Table 1. Expect a mass around 0.2 g. Consider re-doing the trial if your mass is much lower or higher.

Repeat this procedure two more times for a total of 3 trials. You may reuse your hot water bath. You must use the same 125-mL Erlenmeyer flask every time, but make sure you dump out the gas and any condensation before weighing the flask empty. If your aluminum foil gets wet, then you should use a new piece.

Measure the volume of your Erlenmeyer flask using your method, your partner’s method, or a combination of the two. If you need to measure the mass of the water in the flask, be sure to use a high capacity balance such as the CPA324s, which measures up to 320 g. Try not to spill water on the balances. Be sure to clean it up if you do. Record your data in Table 2.

2. Heat of combustion

When it is your turn to use the bomb calorimeter, begin by preparing your sample for analysis:

The experimental sample – this is a clean sample cup with 1.00 g of the volatile liquid inside. Record the exact mass of the volatile liquid in a tared sample cup in Table 3. Cover the sample cup with a piece of parafilm immediately after recording the mass of the sample to ensure none of the sample evaporates.

Your instructor or TA will walk you through measuring the heat of combustion for the sample. You will need the mass of the liquid in order to operate the bomb calorimeter. Record your results in Table 3.

3. Density of liquid state

Use a clean, dry 100 mL volumetric flask with a ground glass stopper for this experiment. How many significant figures do you get? Ask if you are not sure.

Measure the mass of the empty volumetric flask and stopper. Record the data in Table 4.

Fill the volumetric flask exactly to the point where the bottom of the meniscus is on the ground line with the volatile liquid from the bottle labeled “use for density of liquid experiment.” Cover the flask immediately to ensure none of the liquid escapes.

Measure the mass of the filled volumetric flask and stopper. Record the data in Table 4.

4. Boiling point

I will provide a demonstration apparatus for this rather complex set-up, but this is a description of how to build it:

Use the same thermometer that you previously calibrated. If you aren’t sure it is the same one, repeat the calibration procedure and record the offset. Gently push a thermometer through a split rubber stopper. Push the stopper up past the 100 °C mark so you can read temperatures below this mark.

Secure a 5 or 10 mm glass test tube to a thermometer with a rubber band. Be sure not to cover the temperature marks too much; temperatures below 30 °C are probably fine to cover. Make sure the bottom of the thermometer and bottom of the test tube are lined up together.

Use a bunsen clap attached to the rubber stopper to secure the thermometer/test tube apparatus to a ring stand.

Pour about 1 mL of the volatile liquid into the test tube. Drop a boiling chip in the test tube. Then, drop a capillary tube open end down into the test tube, as shown in the illustration. CAUTION! These volatile liquids are all flammable. Keep it away from open flames.

http://www.chem.ucalgary.ca/courses/351/laboratory/boilingpoint.pdf

Below this apparatus, you will need to make a hot water bath. Attach an iron ring to the ring stand about an inch above a bunsen burner, which is connected to the gas with amber tubing. Place a wire gauze on it. Place a 250 mL beaker filled 3/4 of the way with water on the gauze.

Lower the thermometer/test tube apparatus into the water bath, making sure to submerge the volatile liquid completely but allowing sufficient space above the water to ensure no water splashes into the test tube.

Turn on the gas and ignite the bunsen burner. Heat up the water nearly to boiling. Watch the test tube carefully. At first a few bubbles will come out of the capillary tube, but eventually you will see a steady stream of bubbles coming out of the capillary tube. When you see a slow, steady stream of bubbles, turn off the bunsen burner and allow it to cool. Your thermometer should read at least 85 °C before you turn off the flame (that temperature is above all of the possible boiling points). While it is cooling, watch the capillary tube. When you see it start to draw the volatile liquid up into the capillary tube, record the temperature as the boiling point. Be sure to record the temperature to the nearest 0.2 °C. You might miss it if you aren’t watching carefully the whole time it is cooling because it gets slurped up like a straw very quickly. IMPORTANT: be sure to factor in the offset that you found in Part 1 when recording the boiling point temperature. Record data in Table 5.

Repeat for 3 trials total. If three trials do not agree within 3 °C, consult your instructor or TA about performing additional trials. If necessary, refill the test tube and/or hot water bath between trials. Extinguish the flame before refilling or otherwise handling the volatile liquid.

5. Confirmatory test

Once you have completed your calculations and think you know the identity of the unknown, present your answer to the instructor or TA. She will provide a confirmatory test that you can use to validate your answer (or not!). Instructions for your test will be available at that time. Record the name of the test you used and the result. If the first test is negative, you may try a second test.

Calculations

You can now calculate the molar mass of the unknown from the density of the gas. Then, compare the molar mass, the heat of combustion, the boiling point, and the density of the liquid state of the unknown to corresponding data from known substances. Use these results to select the identity of your unknown from this list of possible unknown liquids:

• Molar mass = 78.11 g/mol
• Heat of combustion = 9.992 kcal/g
• Density of liquid = 0.876 g/mL
• Boiling point = 80 °C
• Molar mass = 86.18 g/mol
• Heat of combustion = 0.452 kcal/g
• Density of liquid = 0.659 g/mL
• Boiling point = 68 °C
• Molar mass = 32.04 g/mol
• Heat of combustion = 5.410 kcal/g
• Density of liquid = 0.792 g/mL
• Boiling point = 65 °C
• Molar mass = 46.07 g/mol
• Heat of combustion = 7.086 kcal/g
• Density of liquid = 0.789 g/mL
• Boiling point = 78 °C
• Molar mass = 58.08 g/mol
• Heat of combustion = 7.360 kcal/g
• Density of liquid = 0.791 g/mL
• Boiling point = 56 °C

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Chemistry Predictions 2024 for Leaving Cert Higher Level

Updated December 2023

You may also like: Leaving Cert Chemistry Notes  (€)

Predictions 2024

Experiment 1: Titration

• Titration to determine the concentration of ethanoic acid in vinegar
• Determination of the amount of water of crystallisation in washing crystals
• Iodine-thiosulfate titration

Experiment 2: Organic Chemistry

• Benzoic acid
• Extraction of clove oil from cloves

Experiment 3: Other Experiments

• Measurement of molecular mass of a volatile liquid
• The effect of temperature and concentration on the rate of a reaction
• Estimation of the concentration of free chlorine in swimming-pool water

Question 5:

• Mass spectrometry
• Periodic table e.g. Mendeleev
• Bohr model of the atom
• Flame tests

Question 6: Fuels and heats of reaction

• Thermochemistry calculations
• Fractional distillation of crude oil
• IUPAC names of compounds
• Bomb calorimeter

Question 8: Organic chemistry

• Addition reactions – Ethene and bromine
• Properties of different families of compounds e.g. esters, alcohols, etc.

Questions 7,9:

• Chemical equilibrium: Le Chatelier’s Principle, Calculations, Factors affecting rates of reactions
• Water: Dissolved oxygen, Water and sewage treatment, Hardness of water

Question 10,11:

• Radioactivity
• Acids/Bases and pH calculations
• Stoichiometry
• Balancing redox equations
• Ionic/covalent bonding
• Chromatography
• Q11 (d) will examine options

General patterns of LC Chemistry questions

• For experimental questions, question 1 is a titration, question 2 is based on organics and question 3 is another experiment.
• Determination of the concentration of ethanoic acid in vinegar
• Potassium manganate – ammonium iron (ii) sulfate titration
• Estimation of iron in an iron tablet
• Iodine – thiosulfate titration
• Determination of the percentage of hypochlorite in bleach
• Estimation of the total hardness of a water sample
• Estimation of dissolved oxygen by redox titration
• Standardisation of a hydrochloric acid solution using a standard solution of sodium carbonate
• Recrystallisation and melting point determination of benzoic acid
• Preparation and properties of ethyne
• Preparation of soap
• Preparation and properties of ethene
• Properties of ethanal
• Properties of ethanoic acid
• Preparation of benzoic acid from phenylmethanol
• Separation of a mixture of indicators using chromatography
• Test for anions
• Estimation of the relative molecular mass of a volatile liquid
• Determination of free chlorine in swimming pool water
• Determination of total suspended solids, dissolved solids and pH in a water sample
• Question 4 consists of short questions based on various topics of the course.
• Elements and the periodic table
• Electronic structure of atoms
• Ionic and covalent bonding
• Shapes of molecules and intermolecular forces
• Question 6 is based on fuels and heats of reaction.
• Question 8 is based on organic chemistry.
• The topics of acids and bases, pH, chemical equilibrium and rates of a reaction usually make an appearance somewhere in the long questions. The topics of acids and bases and pH are usually linked in one single long question.
• One of the parts of question 10 will usually be based on organic chemistry or fuels and heats of reaction.
• One of the parts of question 10 or 11 will usually be stoichiometry (calculations based on chemical equations) and another part will usually be based on oxidation and reduction.
• The topics of shapes of molecules, ionic and covalent bonding, intermolecular forces, the atom and radioactivity will usually appear somewhere in the long questions. They are most likely to come up as parts of question 10 or 11.
• Question 11(c) A is based on either atmospheric chemistry or industrial chemistry.
• Question 11(c) B is based on crystals, polymers or metals.

2023: Q1: Potassium manganate-ammonium iron (II) sulfate titration Q2: Preparation of ethyne Q3: Determination of the heat of reaction of hydrochloric acid with sodium hydroxide Q4: Short questions on various topics Q5: Electronic structure and Electronegativity Q6: Fuels and heats of reaction Q7: Acids/Bases/pH and Water Q8: Organic chemistry Q9: Chemical equilibrium Q10: Benzene, Stoichiometry, Radioactivity Q11: Gas laws, Electrolysis, Rates of reactions, Options 2022: Q1: Estimation of the total hardness of a water sample Q2: Ethene, Benzoic acid Q3: Rates of reactions: decomposition of hydrogen peroxide Q4: Short questions on various topics Q5: Electronic structure and Radioactivity Q6: Fuels and heats of reactions Q7: Chemical equilibrium Q8: Organic chemistry Q9: Acids/Bases/pH Q10: Gas laws, Electronegativity, Stoichiometry Q11: Balancing redox equations, Properties of ethanal, Rates of reactions, Options 2021: Q1: Bleach Q2: Soap Q3: Test for anions Q4: Short questions on various topics Q5: History/trends of periodic table

Q6: Fuels and heats of reactions Q7: Acids/Bases/pH Q8: Organic chemistry Q9: Chemical equilibrium Q10: Electron structure, Rates of reactions, Stoichiometry Q11: Ethene, Water, Electron structure/Radioactivity 2020: Q1: Washing crystals Q2: Benzoic acid Q3: Rates of reactions: temperature vs rate Q4: Short questions on various topics Q5: Electronic structure Q6: Fuels and heats of reactions Q7: Acids/Bases/pH Q8: Organic chemistry Q9: Chemical equilibrium Q10: Water, Balancing redox equations, Stoichiometry Q11: Electronegativity, Empirical/Molecular formula, Options

2023: what we predicted

For 2023, it is not mandatory to answer 2 experiment questions as it has been in previous years.

While it is impossible to be certain of what will be on the paper, here are topics to pay special attention to for 2023:

• To prepare a solution of sodium hydroxide and standardise it with a standard hydrochloric acid solution to prepare sodium chloride
• Titration to determine the amount of iron in an iron tablet
• Reactions of ethanal/ethanoic acid
• Preparation of benzoic acid

Experiment 3: Other experiments

• Paper chromatography
• Flame tests/tests for anions
• Answer 8 short questions from 12 on various aspects of the course
• Question 4(l) will be based on Options
• Trends in the periodic table

Question 6 – Fuels and Heats of Reaction

• Heat of combustion/formation calculation
• Bronsted-Lowry vs Arrhenius theory of acids and bases
• pH calculations
• Le Chatelier’s principle

Question 8 – Organic Chemistry

• Properties of different families of compounds e.g. alcohols, aldehydes, carboxylic acids
• Substitution/addition/elimination reactions
• Chemical equilibrium calculations
• Haber/Contact processes

Question 10 – parts a, b, c

• Rates of reaction
• History of the atom

Question 11 – parts a, b, c, d

• Oxidation and reduction
• Part (d) will examine Options

2022: what we predicted

• Preparation and properties of ethyne/ethene
• Determination of the heat of reaction of hydrochloric acid with sodium hydroxide
• Monitor the rate of production of oxygen from hydrogen peroxide using manganese dioxide as a catalyst

Question 7 – Acids and Bases

• Self-ionisation of water

Question 9 – Chemical Equilibrium

2021: what we predicted

• Properties of ethanoic acid/ethanal

Experiment 3: Mixed

• Determination of the concentration of free chlorine in swimming pool water
• Measuring the rate of production of oxygen from hydrogen peroxide
• Halogens as oxidising agents/displacement reactions of metals

Question 4:

• Short questions on various aspects of the course (Options will not be examined)
• History of the atom e.g. discovery of the proton, neutron, electron
• Mass spectrometry, radioactivity

Question 6:

• Fuels and Heats of Reaction – oil refining, petrol, hydrogen as a fuel, heat of combustion

Question 7:

• Acids and bases

Question 8:

• Organic chemistry – properties of different families of compounds, addition/substitution/elimination reactions

Question 9:

• Chemical equilibrium

Questions 10 & 11 (multiple parts):

• Oxidation/reduction reactions
• Chemical bonding – shapes of molecules, electronegativity etc.
• Water – biochemical oxygen demand, wastewater treatment

2020: what we predicted

• Determine the amount of water of crystallisation in hydrated sodium carbonate
• Determine the amount of iron in an iron tablet
• Determine the percentage (w/v) of sodium hypochlorite in bleach 
• Determine the concentration of ethanoic acid in vinegar

Experiment 2: Organic Chemistry 

• Preparation and properties of ethyne and ethene
• To separate a mixture of substances using paper chromatography

Experiment 3: Mix

• Effect of concentration/temperature on the rate of reaction
• Estimate the concentration of free chlorine in swimming-pool water
• Illustration of Le Chatelier’s Principle 
• Halogens as Oxidising Agents/ Displacement Reactions of Metals/ Test for Anions
• Short questions on various aspects of the course

Question 5: 

• Electronic structure of atoms eg. orbitals/ s, p configuration etc. 
• History of the atom eg. Dalton’s Atomic Theory, Discovery of the proton/neutron/electron, Thompson’s plum pudding model, Bohr Model etc. 
• Periodic Table eg. Triads/Octaves/Mendeleev vs Moseley
• Trends in the Periodic Table

Long Questions 6-9: 

• Fuels and Heats of Reaction
• Chemical Equilibrium
• Rates of Reaction

Long Questions with options (Q10 & 11): 

• Stoichiometry 
• Oxidation/Reduction
• Acids and Bases/ pH and Indicators

Options (11c A & B): 

• QA: Atmospheric Chemistry  
• QB: Polymers/Metals

2019: what we predicted

For your reference, this is what we had predicted in 2019:

• Post author: Martina
• Post published: December 20, 2020
• Post category: Chemistry / Predictions

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All Chemistry Definitions to Learn for Leaving Certificate - Higher Level

Shane O'Brien

• Dalton’s Atomic Theory: All matter is made up of indivisible atoms which cannot be created or destroyed.
• William Crookes discovered the Electron. Cathode ray tube.
• JJ Thomson determined that electrons were negative. Found E/M ratio .
• Robert Millikan’s Oil drop experiment discovered size of electron’s charge.
• Plum Pudding Model: Solid positive sphere with electrons embedded.
• Rutherford discovered the Nucleus. Fired alpha particles at gold foil. Proton.
• Chadwick discovered the neutron by firing alpha particles at beryllium.
• Energy level is the fixed energy value that an electron in an atom may have.
• Heisenberg’s Uncertainty Principle states that it is impossible to measure at the same time both the velocity and the position of an electron.
• Orbital is a region of space in which there is a high probability of finding an electron.

Periodic Table

• Element is a substance which cannot be split into simpler substances by chemical means.
• Dobereiner’s Triads are groups of three elements of similar chemical properties in which the atomic weight of the middle element is approximately equal to the average of the other two.
• Newland’s Octaves are groups of elements arranged in order of increasing atomic weight, in which the first and eight element of each group have similar properties.
• Mendeleev’s Periodic Law: When elements are arranged in order of increasing atomic weight, the properties of the elements vary periodically.
• Atomic Number of an atom is the number of protons in the nucleus of that atom.
• Modern Periodic Law: When elements are arranged in order of increasing atomic number, the properties of the elements vary periodically.
• Mass number of an element is the sum of the number of protons and neutrons in the nucleus of an atom of that element.
• Isotopes are atoms of the same element, with the same atomic number, but different mass numbers due to the different number of neutrons in the nucleus.
• Relative Atomic Mass is defined as: The average of the mass numbers of the isotopes of the element, as they occur naturally, taking their abundances into account. Expressed on a scale in which the atoms of the carbon 12 isotope have a mass of exactly 12 units.
• Mass Spectrometer: Vaporisation, Ionisation, Acceleration, Separation, Detection.
• Aufbau Principle states that when building up the electronic configuration of an atom in its ground state, the electrons occupy the lowest available energy level.
• Hund’s Rule of Maximum Multiplicity states that when two or more orbitals of equal energy are available, the electrons occupy them singly before filling them in pairs.
• Pauli Exclusion Principle states that no more than two electrons may occupy an orbital and they must have opposite spin.
• Atomic radius of an atom is defined as half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond.
• First ionisation energy of an atom is the energy required to completely remove the most loosely bound electron from a neutral gaseous atom.

Chemical Bonding

• Compound is a substance that is made up of two or more different elements combined together chemically.
• Octet Rule states that when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell.
• Ion is a charged atom or group of atoms.
• Ionic bond is the force of atttraction between oppositely charged ions in a compound.
• Transition metal is one that forms at least one ion with a partially filled d sublevel.
• Molecule is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.
• Valency of an element is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.
• Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.
• Van der Waals forces are weak attractive forces between molecules resulting from the formation of temporary dipoles.
• Dipole – Dipole forces are forces of attraction between the negative pole of one molecule and the positive pole of another.
• Hydrogen bonds are particular types of dipole-dipole attractions between molecules in which the hydrogen atoms are bonded to nitrogen, oxygen or fluorine. The hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule. Thus, it acts as a bridge between two electronegative atoms.

Chemical Equations – Tests for anions

• Law of Conservation of Mass states that the total mass of the products of a chemical reaction is the same as the total mass of the reactants.
• Law of Conservation of Matter states that in any chemical reaction, matter is neither created nor destroyed, but merely changes from one form into another.

Radioactivity

• Radioactivity is the spontaneous breaking up of unstable nuclei with the emission of one or more types of radiation.
• Half Life of an element is the time taken for half of the nuclei in any given sample to decay.

Mole Concept

• One mole of a substance is the amount of that substance that contains 6 x 10^23 particles of that substance. Mass = Rel.Molecular Mass x Number of moles
• Relative Molecular Mass of a molecule is the sum of the relative atomic masses of all the atoms in a molecule of the compound.

Properties of Gases

• A Gas is defined as a substance that has no well-defined boundaries but diffuses rapidly to fill any container in which it is placed.
• Standard Temperature = 273K
• Boyle’s Law states that, at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure.
• Charles’ Law states that, at constant pressure, the volume of a fixed mass of gas is directly proportional to its temperature measured on theKelvin scale.
• Gay-Lussac’s Law of Combining Volumes:  In a reaction between gases, the volumes of the reacting gases and the volumes of any gaseous products are in the ratio of small whole numbers provided the volumes are measured at the same temperature and pressure.
• Avagadro’s Law states that equal volumes of gases contain equal numbers of molecules, under the same conditions of pressure and temperature.
• Ideal Gas is one that obeys all the assumptions of the kinetic theory of gases under all conditions of temperature and pressure.
• Real gases vs Ideal Gases:  Forces of attraction/repulsion between the molecules. Volume of molecules is not negligible.

Acids & Bases, pH

• Acid: Substance that dissociates in water to produce H+ ions.
• Base: Substance that dissociates in water to produce OH-ions.
• Brønsted Lowry definitions:  Acid is a proton donor. Base is a proton acceptor.
• Conjugate Acid: Acid will change to a conjugate base when it donates a proton.
• Conjugate Base: Base will change to a conjugate acid when it accepts a proton. The stronger the acid, the weaker the conjugate base.
• Conjugate Acid-Base Pair is any pair consisting of an acid and a base which differ by one proton.
• Neutralisation is the reaction between an acid and a base to form a salt and water.

Volumetric Analysis

• Concentration of a solution is the amount of solute that is dissolved ina given volume of solution.
• Primary standard is a substance which can be obtained in a stable, pure and soluble solid form so that it can be weighed out and dissolved in water to give a solution of accurately known concentration.

Oxidation & Reduction

• OIL RIG -Oxidation is loss of electrons, Reduction is gain of electrons. 
• Oxidising agent is a substance that brings about oxidation in other substances.
• Reducing agent is a substance that brings about reduction in other substances.
• Oxidation number is the charge that an atom has or appears to have when electrons are distributed according to certain rules.
• Oxidation = Increase in ON
• Reduction = Decrease in ON

Rates of Reaction

• Rate of Reaction is defined as the change in concentration per unit time of any one reactant or product.
• Catalyst is a substance that alters the rate of chemical reactio but is not used up in the reaction.
• Activation energy is the minimum energy which colliding particles must have for a reaction to occur.

Chemical Equilibrium

• Chemical Equilibrium is a state of dynamic balance where the rate of the forward reaction equals the rate of the reverse reaction.
• Le Chatelier’s Principle: If a stress is applied to a system at equilibrium, the system readjusts to relieve the stress applied.
• Hard water is water that will not easily form a lather with soap. Hardness in water is caused by the presence of Ca2+ or Mg2+ ions.
• Water Treatment:  Screening – Flocculation – Sedimentation – Filtration – Chlorination – Flouridation – pH adjustment
• Flocculant is a chemical added to water to coagulate suspended particles.
• B.O.D -Biochemical Oxygen demand is defined as the amount of dissolved oxygen consumed by bilogical action when a sample of water is kept at 20 degrees in the dark for 5 days.
• Eutrophication is the enrichment of water with nutrients, which leads to excessive growth of algae.

Electrochemistry

• Electrolysis is the use of electricity to bring about a chemical reaction. 
• Electroplating is a process where electrolysis is used to put a layer of one metal on the surface of another.
• The electrochemical series is a list of elements in order of their standard electrode potentials.

Organic Chemistry

• Hydrocarbon is a compound that contains only hydrogen and carbon. 
• Homologous Series is defines as a series of compounds of uniform chemical type, showing gradations in physical properties, having a general formula for its members, each member having a similar method of preparation, each member differing from the previous by a CH2 unit.
• Structural isomers are compounds with the same molecular formula but different structural formulae.
• Aliphatic compound is an organic compound that consists of open chains of carbon atoms and closed chain compounds (rings) that resemble them in chemical properties.
• Aromatic compounds are compounds that contain a benzene ring in their structure.
• Octane Number of a fuel is a measure of its tendency to resist knocking.
• Catalytic Cracking is the breaking down of long-chain hydrocarbon molecules into short chain molecules for which there is greater demand.
• Heat of Reaction is the heat change when the number of moles of reactants indicated in the balanced equation for the reaction react completely.
• Heat of Combustion of a substance is the heat change when one mole of the substance is burned completely in excess oxygen.
• Kilogram Calorific Value of a a fuel is the heat energy produced when one kg of the fuel is burned completely in oxygen.
• Bond energy is the energy required to break one mole of covalent bonds and to separate the neutral atoms completely from each other.
• Heat of Neutralisation is the heat change when one mole of H+ ions from an acid reacts with one mole of OHions from a base.
• Hess’s Law states that if a chemical reaction takes place in a number of stages, the sum of the heat changes in the separate stages is equal to the heat change if the reaction is carried out in one stage.
• Law of Conservation of Energy states that energy cannot be created or destroyed, but can only be converted from one form of energy into another.
• Functional group is an atom or group of atoms which is responsible for the characteristic properties of a series of organic compounds.
• Substitution reaction is a chemical reaction in which an atom or group of atoms in a molecule is replaced by another atom or group of atoms.
• Mechanism of a reaction is the detailed step by step description of how the overall reaction occurs.
• Addition reaction is one in which two substances react together to form a single substance.
• Polymers are long chain molecules made by joining together many small chain molecules.
• Organic Synthesis is the process of making organic compounds from simpler starting materials.
• Chromatography is a separation technique in which a mobile phase carrying a mixture moves in contact with a selectively adsorbent stationary phase.
• Reflux is literally a process in which the reactants are boiled in a container usually a round bottom flask. The vapour that is formed in the boiling process remains in the container (e.g. Round bottom flask) because of the cooking process involving the condenser attached to the round bottom flask. It helps to keeps the product formed in the container without escaping. The product that is in the vapour state returns to it’s liquid form.