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3bf487dc-3f9b-45fc-8132-84460b422a56 < Back Previous Next Time, temperature, mass & volume Time Time can be measured using a stopwatch or stopclock which are usually accurate to one or two decimal places The units of time normally used are seconds or minutes Other units may be used for extremely slow reactions (e.g. rusting) Remember: 1 minute = 60 seconds Examiner Tips and Tricks Careful: Units of time often cause issues in results tables. If the display on a stopwatch showed 1:30. The incorrect time to record would be 1.30 minutes. The correct time would be 1.5 minutes. To avoid any confusion, if the time intervals are less than a minute, it is best / easire to change the recorded units to seconds. So, the same stopwatch display would be recorded as 90 seconds. Temperature Temperature is measured with a thermometer or digital temperature probe Laboratory thermometers usually have a precision of a half or one degree Digital temperature probes are available which are more precise than traditional thermometers and can often read to 0.1 oC Traditional thermometers rely upon the uniform expansion and contraction of a liquid substance with temperature Digital temperature probes can be just as, if not, more accurate than traditional thermometers The units of temperature are degrees Celsius (ºC) Mass Mass is measured using a digital balance which normally gives readings to two decimal places Balances should be tared (set to zero) before use Balances should also be allowed time to settle on a final measurement / reading before it is recorded The standard unit of mass in kilograms (kg) However, in chemistry grams (g) are most often used Remember: 1 kilogram = 1000 grams Volumes of liquid The volume of a liquid can be determined using different pieces of apparatus The choice of apparatus depends on the level of accuracy needed Three common pieces of apparatus for measuring the volume of a liquid are: Burettes Volumetric pipettes Measuring cylinders Burettes are the most accurate way of measuring a variable volume of liquid between 0 cm3 and 50 cm3 They are most commonly used in titrations Careful: Read the burette scale from top to bottom as 0.00 cm3 is at the top of the column Volumetric pipettes are the most accurate way of measuring a fixed volume of liquid, They have a scratch mark on the neck which is matched to the bottom of the meniscus to make the measurement A pipette filler is used to draw the liquid into the volumetric pipette The most common volumes for volumetric pipettes are 10 cm3 and 25 cm3 Measuring cylinders are used when approximate volumes are required (accuracy is not an important factor) These are graduated (have a scale so can be used to measure) Measuring cylinders typically range from 10 cm3 to 1 litre (1 dm3) Whichever apparatus you use, you may see markings in millilitres, ml, which are the same as a cm3 Volumes of gas For some experiments, the volume of a gas produced needs to be measured This is typically done by using one of the following methods: Using a gas syringe By downward displacement of water A gas syringe is more precise and accurate than downward displacement of water Diagram of the set-up for an experiment involving a gas syringe Downward displacement of water is where a measuring cylinder is inverted in water to collect the gas produced This method does not work if the gas is soluble in water Diagram of the set-up for an experiment collecting gas by downward displacement of water If the gas happens to be heavier than air and is coloured, the cylinder does not need to be inverted Advantages & disadvantages of methods & apparatus In the lab, we often have choices of different apparatus to do the same job Evaluating which piece of apparatus is the best one to use is part of good experimental planning and design This means appreciating some of the advantages and disadvantages of laboratory apparatus Advantages and disadvantages of lab apparatus Apparatus Advantage Disadvantage Temperature probe More precise readings Easy to make multiple repeat readings Can be automated to run over long periods of time Can be corroded by some reagents More expensive (to replace) Volumetric pipette Accurate measurement of a fixed volume Harder to use than a normal pipette Only measures one fixed volume Gas syringe Easy to set up Keeps the gas dry The syringe can stick Collects limited volumes Expensive and delicate / fragile Microscale experiments Less wasteful Saves energy Safer Hard to see what's happening Lose a lot of material separating / purifying the products Five pieces of apparatus that can be used to measure the volume of a liquid. They all have their pros and cons Planning your method Good experimental design includes the answers to questions like Have I chosen a suitable apparatus for what I need to measure? Is it going to give me results in an appropriate time frame? Is it going to give me enough results to process, analyse and make conclusions? Does it allow for repetitions to check how reliable my results are? Does my plan give a suitable range of results? How can I be sure my results are accurate ? Have I chosen an appropriate scale of quantities without being wasteful or unsafe? You may be asked about experimental methods in exam questions and your experience and knowledge of practical techniques in chemistry should help you to spot mistakes and suggest improvements Solutions You need to know all the following terms used when describing solutions: Terminology about solutions table Term Meaning Example Solvent The liquid in which a solute dissolves The water in seawater Solute The substance which dissolves in a liquid to form a solution The salt in seawater Solution The mixture formed when a solute is dissolved in a solvent Seawater Saturated solution A solution with the maximum concentration of solute dissolved in the solvent Seawater in the Dead Sea Soluble A substance that will dissolve Salt is soluble in water Insoluble A substance that will not dissolve Sand is insoluble in water Filtrate The liquid or solution that has passed through a filter Fresh coffee in a cup Residue The substance that remains after evaporation, distillation, filtration or any other similar process Coffee grounds in filter paper Acid-base titrations Titrations are a method of analysing the concentration of solutions They can determine exactly how much alkali is needed to neutralise a quantity of acid – and vice versa You may be asked to perform titration calculations to determine the moles present in a given amount or the concentration / volume required to neutralise an acid or a base Titrations can also be used to prepare salts Apparatus 25 cm3 volumetric pipette Pipette filler 50 cm3 burette 250 cm3 conical flask Small funnel 0.1 mol / dm3 sodium hydroxide solution Sulfuric acid of unknown concentration A suitable indicator Clamp stand, clamp & white tile The steps in performing a titration Method Use the pipette and pipette filler and place exactly 25 cm3 sodium hydroxide solution into the conical flask Using the funnel, fill the burette with hydrochloric acid placing an empty beaker underneath the tap. Run a small portion of acid through the burette to remove any air bubbles Record the starting point on the burette to the nearest 0.05 cm3 Place the conical flask on a white tile so the tip of the burette is inside the flask Add a few drops of a suitable indicator to the solution in the conical flask Perform a rough titration by taking the burette reading and running in the solution in 1 – 3 cm3 portions, while swirling the flask vigorously Quickly close the tap when the end-point is reached The endpoint is when one drop causes a sharp colour change Record the volume of hydrochloric acid added, in a suitable results table as shown below Make sure your eye is level with the meniscus Repeat the titration with a fresh batch of sodium hydroxide As the rough end-point volume is approached, add the solution from the burette one drop at a time until the indicator just changes colour Record the volume to the nearest 0.05 cm3 Repeat until you achieve two concordant results (two results that are within 0.1 cm3 of each other) to increase accuracy Rough titre Titre 1 Titre 2 Titre 3 Final reading (cm3) First reading (cm3) Titre (cm3) Examiner Tips and Tricks Common errors during a titration include: Not removing the funnel from the burette This can lead to some liquid dripping into the burette and cause false / high readings Not filling the jet space of the burette The jet space is the part of the burette after the tap Not filling this space can lead to false readings Reading the volume from the burette incorrectly Readings should be taken from the bottom of the meniscus Careful: The scale on the burette has 0.0 cm3 at the top and 50 cm3 (typically) at the bottom Indicators Indicators are used to show the endpoint in a titration Wide range indicators such as litmus are not suitable for titration as they do not give a sharp colour change at the endpoint However, methyl orange and phenolphthalein are very suitable Some of the most common indicators with their corresponding colours are shown below: Common acid-base indicators Indicator Colour in acid Colour in alkali Colour in neutral Litmus solution Red Blue Purple Red litmus paper Stays red Turns blue No change Blue litmus paper Turns red Stays blue No change Methyl orange Red Yellow Orange Phenolphthalein Colourless Pink Colourless Thymolphthalein Colourless Blue Colourless Paper chromatography Chromatography is used to separate substances and provide information to help identify them The components have different solubilities in a given solvent E.g. Different coloured inks that have been mixed to make black ink A pencil line is drawn on chromatography paper and spots of the sample are placed on it A pencil is used for this as ink would run into the chromatogram along with the samples The paper is then lowered into the solvent container, making sure that the pencil line sits above the level of the solvent so the samples don’t wash into the solvent container The solvent used is usually water but it can be other substances such as ethanol The solvent travels up the paper by capillary action , taking some of the coloured substances with it Different substances have different solubilities so they will travel at different rates, causing the substances to spread apart Those substances with higher solubility will travel further than the others How to carry out chromatography The pigments in ink can be analysed using paper chromatography Interpret simple chromatograms We can use a chromatogram to compare the substances present in a mixture to known substances and make assumptions Pure substances will produce only one spot on the chromatogram Impure substances will produce more than one spot on the chromatogram If two or more substances are the same, they will produce identical chromatograms If the substance is a mixture , it will separate on the paper to show all the different components as separate spots It is common practice to include a known compound as a reference spot This can help match up to an unknown spot or set of spots in order to identify it Example chromatogram results The brown ink has separated showing a spot of red ink, blue ink and yellow ink We can draw several conclusions from this chromatogram: The brown ink is a mixture as there are three dots Red, yellow and blue are pure as there is only one dot for each The brown ink contains red, blue and yellow as the dots are in line with one another horizontally Examiner Tips and Tricks Chromatograms in exams will be in black and white so to identify whether a mixture contains a known sample, the dots need to be in line with one another. Locating agents Extended tier only For chromatography to be useful, the chemist needs to be able to see the components move up the paper This is not the case for colourless substances such as amino acids or sugars Locating agents can be used to see the spots These are substances which react with the sample and produce a visible / coloured spot for the product(s) The chromatogram is treated with the agent after the chromatography run has been carried out, making the sample runs visible to the naked eye Retention factor (Rf) values Extended tier only R f values are used to identify the components of mixtures The R f value of a particular compound is always the same However, it does depend on the solvent used If the solvent is changed then the R f value changes Calculating the R f value allows chemists to identify unknown substances because it can be compared with the R f values of known substances under the same conditions The retention factor, R f, is calculated by the equation: R f = The R f value: Is a ratio Has no units Will always be less than 1 Worked Example A student obtained the following chromatogram when carrying out chromatography. Calculate the R f value of the substance. Answer: The R f value of the substances in the chromatogram above can be calculated by: R f = = = 0.5 Examiner Tips and Tricks When you calculate R f values in exams, make sure to use your ruler carefully to measure the distance moved by the solvent and the substance as mark schemes can be strict about the values accepted for these. Filtration & crystallisation The choice of separation technique depends on the substances being separated All techniques rely on a difference in properties of the chemicals in the mixture This is usually a physical property such as boiling point Separating a mixture of solids Differences in solubility can be used to separate solids For a difference in solubility, a suitable solvent must be carefully chosen Only the desired substance should dissolve in the solvent Other substances or impurities in the mixture should not dissolve in the solvent For example, to separate a mixture of sand and salt: Water is a suitable solvent because salt is soluble in water, but sand is insoluble in water Filtration This technique is used to separate an undissolved solid from a mixture of the solid and a liquid / solution ( e.g. sand from a mixture of sand and water) Centrifugation can also be used for this mixture A filter paper is placed in a filter funnel above another beaker The mixture of insoluble solid and liquid is poured into the filter funnel The filter paper will only allow small liquid particles to pass through in the filtrate Solid particles are too large to pass through the filter paper so will stay behind as a residue Filtration of a mixture of sand and water Crystallisation This method is used to separate a dissolved solid from a solution A simple application of this is to heat a solution to boiling, remove the heat and leave the solvent to evaporate A more common application of this is sometimes called crystallisation This is when the solid is more soluble in hot solvent than in cold, e.g. copper sulphate from a solution of copper(II) sulphate The solution is heated, allowing the solvent to evaporate and leaving a saturated solution behind You can test if the solution is saturated by dipping a clean, dry, cold glass rod into the solution If the solution is saturated, crystals will form on the glass rod when it is removed and allowed to cool The saturated solution is allowed to cool slowly Solids will come out of the solution as the solubility decreases This will be seen as crystals growing The crystals are collected by filtration They are then washed with distilled water to remove any impurities Finally, they are allowed to dry Common places to dry crystals are between sheets of filter paper or in a drying oven The process of crystallisation The solution is slowly heated to remove around half of the liquid. The remaining liquid will evaporate slowly Examiner Tips and Tricks In exams, you need to be specific that no more than half of the solution is removed by direct heating or you may lose a mark. Distillation: simple & fractional Simple distillation Distillation is used to separate a liquid and soluble solid from a solution (e.g. water from a solution of saltwater) or a pure liquid from a mixture of liquids The solution is heated and pure water evaporates producing a vapour which rises through the neck of the round-bottomed flask The vapour passes through the condenser , where it cools and condenses, turning into pure water which is collected in a beaker After all the water is evaporated from the solution, only the solid solute will be left behind Simple distillation apparatus Diagram showing the distillation of a mixture of salt and water Simple distillation can be used to separate the products of fermentation, such as alcohol and water However, fractional distillation is a more effective separation technique, commonly used when the boiling points of the liquids are close and/or a higher degree of purity is required, such as crude oil Fractional distillation Used to separate two or more liquids that are miscible with one another (e.g. ethanol and water from a mixture of the two) The solution is heated to the temperature of the substance with the lowest boiling point This substance will rise and evaporate first The vapours will pass through a condenser, where they cool and condense The condensed liquid is then collected in a beaker All of the substance is evaporated and collected, leaving behind the other component(s) of the mixture For water and ethanol: Ethanol has a boiling point of 78 ºC Water has a boiling point of of 100 ºC The mixture is heated until it reaches 78 ºC, at which point the ethanol distills out of the mixture and into the beaker When the temperature starts to increase to 100 ºC heating should be stopped as the water and ethanol are now separated Fractional distillation of a mixture of ethanol and water An electric heater is safer to use when there are flammable liquids present The separation of the components in petroleum is achieved by fractional distillation on an industrial scale Fractional distillation of crude oil is not carried out in school laboratories due to the toxic nature of some of the components of the crude oil, but it can sometimes be simulated using a synthetic crude oil made specially for the demonstration Worked Example A student is given a mixture of calcium sulfate, magnesium chloride and water. The table below shows some information about calcium sulfate and magnesium chloride. substance solubility in water state at room temperature calcium sulfate insoluble solid magnesium chloride soluble solid How does the student obtain magnesium chloride crystals from the mixture? Crystallisation followed by distillation Crystallisation followed by filtration Distillation followed by crystallisation Filtration followed by crystallisation Answer The correct answer is D because: The difference in solubility in water means the first step is to make a solution The magnesium chloride will dissolve, but the solid calcium sulfate will be left behind The mixture is filtered to remove the calcium sulfate and then evaporated and crystallised to obtain magnesium chloride crystals Examiner Tips and Tricks You may be asked how to separate a mixture of gases: One method involves cooling the gaseous mixture sufficiently to liquefy all of the gases, which are then separated by fractional distillation. They can also be separated by diffusion, where the boiling points are very close or it is impractical or expensive to use fractional distillation. Assessing purity Pure substances melt and boil at specific and sharp temperatures For example, water has a boiling point of 100°C and a melting point of 0°C Mixtures have a range of melting and boiling points as they consist of different substances that melt or boil at different temperatures Therefore, melting and boiling point data can be used to distinguish pure substances from mixtures An unknown pure substance can be identified by experimentally determining its melting point and boiling point and comparing them to literature values / data tables Boiling points are commonly determined by distillation Melting point analysis is routinely used to assess the purity of drugs for example This is done using a melting point apparatus which allows you to slowly heat up a small amount of the sample, making it easier to observe the exact melting point Melting point test using an oil bath This is then compared to data tables The closer the measured value is to the actual melting or boiling point, the purer the sample is If the sample contains impurities: The boiling point may appear higher than the sample's actual boiling point The melting point may appear lower than the sample's actual melting point Identification of anions Negatively charged non-metal ions are known as anions You must be able to describe the tests for the following ions: Carbonate ions, CO32– Halide ions, Cl– , Br– , I– Nitrate ions, NO3– Sulfate ions, SO42– Sulfite ions, SO32– Test for carbonate ions Carbonate compounds contain the carbonate ion, CO32- The test for the carbonate ion is: Add dilute acid Bubble the gas released through limewater Limewater turns cloudy if the carbonate ion is present If a carbonate compound is present then fizzing / effervescence should be seen as CO2 gas is produced, which forms a white precipitate of calcium carbonate when bubbled through limewater: CO32- (aq) + 2H+ (aq) → CO2 (g) + H2O (l) CO2 (g) + Ca(OH)2 (aq) → CaCO3(s) + H2O(l) The white precipitate turns limewater cloudy Testing for carbonate ions Limewater turns milky in the presence of carbon dixoide caused by the formation of insoluble calcium carbonate Examiner Tips and Tricks If you are asked to describe the test for carbonate ions, make sure that you say: Bubble the gas produced through limewater, which turns cloudy if the carbonate ion is present Just saying that limewater turns cloudy is not enough This isn't describing the test, it is stating the result Test for halide ions Halide ions are the negative ions / anions formed by the elements in Group 7 The test for the halide ions is: Acidify the sample with nitric acid Add silver nitrate solution, AgNO3, A silver halide precipitate forms if a halide ion is present The precipitate is indicated by the state symbol (s) The colour of the silver halide precipitate depends on the halide ion: The chloride ion forms a white precipitate of silver chloride potassium chloride + silver nitrate → potassium nitrate + silver chloride KCl (aq) + AgNO3 (aq) → KNO3 (aq) + AgCl (s) The bromide ion forms a cream precipitate of silver bromide potassium bromide + silver nitrate → potassium nitrate + silver bromide KBr (aq) + AgNO3 (aq) → KNO3 (aq) + AgBr (s) The iodide ions forms a yellow precipitate of silver iodide potassium iodide + silver nitrate → potassium nitrate + silver iodide KI (aq) + AgNO3 (aq) → KNO3 (aq) + AgI (s) Testing for halide ions Each silver halide produces a precipitate of a different colour Examiner Tips and Tricks The acidification step in the halide ion test must be done with nitric acid rather than hydrochloric acid. HCl contains the chloride ion which would interfere with the results. Test for nitrate ions Nitrate compounds contain the nitrate ion, NO3– The test for the nitrate ion is Add aqueous NaOH and aluminium foil Warm gently and test the gas released The gas given off is ammonia, NH3 Ammonia is a gas with a characteristic sharp choking smell that turns damp red litmus paper blue Test for sulfate ions Sulfate compounds contain the sulfate ion, SO42- The test for the sulfate ion is: Acidify the sample with dilute nitric acid Add a few drops of barium nitrate solution A white precipitate of barium sulfate is formed, if the sulfate ion is present Ba2+ (aq) + SO42- (aq) → BaSO4 (s) The test can also be carried out with barium nitrate solution Testing for sulfate ions A white precipitate of barium sulfate is a positive result for the presence of sulfate ions Examiner Tips and Tricks Nitric is added first to remove any carbonates which may be present which would also produce a precipitate and interfere with the results. Test for sulfite ions Sulfite compounds contain the sulfite ion, SO32- The test for the sulfite ion is: Add dilute acid Warm the mixture gently Bubble the gas released through potassium manganate(VII) solution The potassium manganate(VII) solution changes from purple to colourless if the sulfite ion is present Examiner Tips and Tricks For qualitative inorganic analysis, there will be one test for the metal cation and another test for the non-metal anion . If you are an extended level student you may be asked to write balanced ionic equations for cation and anions tests, so make sure you know the formulae of all the ions and precipitates formed. Identification of cations Test for ammonium ions Ammonium ions, NH4+, can be identified by gently warming a solution containing the ions with sodium hydroxide solution The sodium hydroxide solution is a source of hydroxide ions, OH–, for the test This releases ammonia gas which turns damp red litmus paper blue Testing for ammonium ions Heating ammonium ions with sodium hydroxide solution releases ammonia gas which turns damp red litmus blue Metal cations in aqueous solution can be identified by the colour of the precipitate they form on addition of sodium hydroxide and ammonia Most transition metals produce hydroxides with distinctive colours Test for metal ions with sodium hydroxide solution If a small amount of sodium hydroxide solution is used, the resulting metal hydroxide normally precipitates out of solution If excess sodium hydroxide solutionis used, some of the precipitates may re-dissolve For this reason, just a few drops of sodium hydroxide solutionare added at first and very slowly The sodium hydroxide test for the metal ion is: Add a few drops of sodium hydroxide solution Record any colour changes or precipitates formed Add excess sodium hydroxide solution Record any colour changes or changes to precipitates Test for metal ions with ammonia solution If a small amount of ammonia solution is used, the resulting metal hydroxide normally precipitates out of solution If excess ammonia solution is used, some of the precipitates may re-dissolve For this reason, just a few drops of ammonia solution are added at first and very slowly The ammonia test for the metal ion is: Add a few drops of ammonia solution Record any colour changes or precipitates formed Add excess ammonia solution Record any colour changes or changes to precipitates Metal ion tests summary Initially, sodium hydroxide solution and ammonia solution give the same results for 2 - 3 drops This is because they both contain the hydroxide ion, OH– Metal Ion Addition of 2-3 drops of NaOH or ammonia Addition of excess NaOH Addition of excess ammonia Al3+ White precipitate forms Precipitate dissolves to form a colourless solution Precipitate does not dissolve Ca2+ White precipitate forms Precipitate does not dissolve Precipitate does not dissolve Cr3+ Green precipitate forms Precipitate dissolves to form a green solution Precipitate does not dissolve Cu2+ Light blue precipitate forms Precipitate does not dissolve Precipitate dissolves to form a dark blue solution Fe2+ Green precipitate forms Precipitate does not dissolve Precipitate does not dissolve Fe3+ Brown precipitate forms Precipitate does not dissolve Precipitate does not dissolve Zn2+ White precipitate forms Precipitate dissolves to form a colourless solution Precipitate dissolves to form a colourless solution Analysing results The tables above contain the results for all metal cations included in the syllabus If a precipitate is formed from either sodium hydroxide or ammonia solution, then the hydroxide is insoluble in water For example, zinc chloride: ZnCl2 (aq) + 2NaOH (aq) → Zn(OH)2 (s) + 2NaCl (aq) There are 3 metal ions that all form white precipitates: Aluminium ions, Al3+ Calcium ions, Ca2+ Zinc ions, Zn2+ Calcium ions, Ca2+, can be easily distinguished from Zn2+ and Al3+ The white precipitate of calcium hydroxide does not dissolve in excess sodium hydroxide solution The white precipitates of zinc hydroxide and aluminium hydroxide dissolve in excess sodium hydroxide solution Zinc ions, Zn2+, can then be distinguished from Al3+ ions as The white precipitate of zinc hydroxide dissolves in excess ammonia solution The white precipitate of aluminium hydroxide does not dissolve in excess ammonia solution Examiner Tips and Tricks The ammonia or sodium hydroxide solution must be added very slowly. If it is added too quickly and the precipitate is soluble in excess, then you run the risk of missing the formation of the initial precipitate, which dissolves as quickly as it forms if excess solution is added. Be sure to distinguish between the term “ colourless ” and “ clear ”. A solution that loses its colour has become colourless. A clear solution is one that you can see through such as water. Solutions can be clear and have colour eg. dilute copper sulphate. Flame tests for metal ions The flame test is used to identify the metal cations by the colour of the flame they produce Ions from different metals produce different colours Dip the loop of an unreactive metal wire such as nichrome or platinum in concentrated acid and then hold it in the blue flame of a Bunsen burner until there is no colour change This is an important step as the test will only work if there is just one type of ion present Two or more ions means the colours will mix, making identification erroneous This cleans the wire loop and avoids contamination A small sample of the compound is placed on an unreactive metal wire loop such as nichrome or platinum Dip the loop into the solid sample / solution and place it in the edge of the blue Bunsen flame Avoid letting the wire get so hot that it glows red otherwise this can be confused with a flame colour Diagram showing the technique for carrying out a flame test The colour of the flame is observed and used to identify the metal ion present: Cation Flame Colour Li+ Crimson Na+ Yellow K+ Lilac Ca2+ Red Ba2+ Apple-green Cu2+ Blue-green Metal ions form distinctive coloured flames Experimental Techniques Next Topic

< Back Chapter 10 prerequisite Previous Next 🌈🌟📘 Prerequisites for Chapter 10: Periodicity 📘🌟🌈Before diving into 🚀 Chapter 10 , which deals with Periodicity , students must have a solid understanding of the following concepts:🔬 1. Basic Atomic Structure 🧪Understand protons, neutrons, and electrons.🔬 2. The Periodic Table 📊Be familiar with the layout of the periodic table and the properties of elements based on their position.🔬 3. Electron Configuration 🌀Understand how electrons are arranged in atoms.🔬 4. Trends in the Periodic Table 📈Understand the trends in atomic size, ionization energy, electronegativity, and metallic character.🌟 20 Multiple Choice Questions for Chapter 10: Periodicity 🌟What is the term for the repeating pattern of chemical properties in elements in the periodic table? a) Periodicity b) Atomicity c) Reactivity d) IsotopyAs you move from left to right across a period, what generally happens to the atomic size? a) Increases b) Decreases c) Remains the same d) Increases then decreasesWhat is the energy required to remove an electron from an atom called? a) Electron affinity b) Ionization energy c) Electronegativity d) Atomic radiusWhich group of elements is known for being unreactive? a) Alkali metals b) Alkaline earth metals c) Halogens d) Noble gasesWhat is the term for the ability of an atom to attract electrons in a chemical bond? a) Electron affinity b) Ionization energy c) Electronegativity d) Atomic radiusWhich element has the highest electronegativity? a) Fluorine b) Oxygen c) Nitrogen d) ChlorineAs you move down a group in the periodic table, what generally happens to the atomic size? a) Increases b) Decreases c) Remains the same d) Increases then decreasesWhat is the term for the half the distance between the nuclei of two bonded atoms of the same element? a) Electron affinity b) Ionization energy c) Electronegativity d) Atomic radiusWhich group of elements is highly reactive and has one electron in their outermost energy level? a) Alkali metals b) Alkaline earth metals c) Halogens d) Noble gasesWhat is the general trend in ionization energy as you move from left to right across a period? a) Increases b) Decreases c) Remains the same d) Increases then decreasesWhich element is located in Group 2 and Period 3 of the periodic table? a) Magnesium b) Calcium c) Sodium d) AluminumWhat is the general trend in electronegativity as you move down a group in the periodic table? a) Increases b) Decreases c) Remains the same d) Increases then decreasesWhich element is known as the 'King of Chemicals' due to its high reactivity? a) Oxygen b) Fluorine c) Chlorine d) HydrogenWhat is the electron configuration of an atom in the noble gas group? a) Fully filled s and p subshells b) Half-filled s subshell c) Fully filled d subshell d) Half-filled p subshellWhich element has the lowest ionization energy? a) Helium b) Francium c) Fluorine d) CesiumWhat is the general trend in metallic character as you move from left to right across a period? a) Increases b) Decreases c) Remains the same d) Increases then decreasesWhich element is a liquid at room temperature and is located in Group 17 of the periodic table? a) Bromine b) Iodine c) Fluorine d) ChlorineWhat is the term for the energy change when an electron is added to an atom? a) Electron affinity b) Ionization energy c) Electronegativity d) Atomic radiusWhich of the following elements is a metalloid? a) Silicon b) Sodium c) Sulfur d) SilverWhat is the general trend in atomic radius as you move down a group in the periodic table? a) Increases b) Decreases c) Remains the same d) Increases then decreases🌈🌟 Answers 🌟🌈a) Periodicityb) Decreasesb) Ionization energyd) Noble gasesc) Electronegativitya) Fluorinea) Increasesd) Atomic radiusa) Alkali metalsa) Increasesa) Magnesiumb) Decreasesb) Fluorinea) Fully filled s and p subshellsb) Franciumb) Decreasesa) Brominea) Electron affinitya) Silicona) Increases

Previous All Content Next Chapter 5 SABIS Grade 10 Lesson 1 Intro XXXX Hello students! Today, we are going to explore some basic concepts related to gases. Gases are one of the three classical states of matter and understanding their behavior is fundamental in chemistry. We will break down these concepts into simple terms, provide examples, and have small questions to check your understanding. Let's get started! Part 1: Molar Volume of Gases Concept 1: Molar Volume in Different States Molar volume is the volume occupied by one mole of a substance. For gases, the molar volume is much larger compared to the liquid state. In fact, it is about 1000 times larger! 🔎 Example: Imagine a balloon filled with water and another filled with air. The air-filled balloon can expand much more than the water-filled balloon for the same amount of substance. 📝 Quick Question: Why do you think gases have a larger molar volume compared to liquids? 🌟 Answer : Gases have a larger molar volume compared to liquids because the particles in a gas are much more spread out than in a liquid. In a gas, the molecules are in constant motion and are far apart from each other, with a lot of empty space between them. This allows gases to expand and fill the volume of their container. In contrast, in a liquid, the molecules are much closer together and have less freedom to move around, which results in a smaller molar volume. Additionally, the attractive forces between molecules in a liquid are stronger than in a gas, keeping the molecules in close proximity to each other. Question 1: Which of the following best explains why gases have a larger molar volume than liquids? a) Gases have stronger intermolecular forces than liquids. b) The particles in a gas are more spread out and have more empty space between them compared to a liquid. c) Gases are always at a higher temperature than liquids. d) The particles in a gas are larger than those in a liquid. Correct Answer: b) The particles in a gas are more spread out and have more empty space between them compared to a liquid. Question 2: One mole of any gas at Standard Temperature and Pressure (STP) occupies a volume of: a) 1 liter b) 22.4 liters c) 6.022 x 10²³ liters d) 1000 liters Correct Answer: b) 22.4 liters Question 3: If you have equal moles of helium gas and water in their respective containers, which one will occupy a larger volume? a) Helium gas b) Water c) Both will occupy the same volume d) Cannot be determined Correct Answer: a) Helium gas Concept 2: Volume of a Gas The volume of a gas is equal to the volume of the container it occupies. Gases have the ability to fill their container regardless of the shape or size. 🔎 Example: If you have a gas in a small bottle and you release it into a big room, the gas will spread out and fill the entire room. 📝 Quick Question: If you have a balloon filled with helium and you release the helium into a classroom, will the helium occupy the entire classroom? Why or why not? 🌟 Answer : Yes, the helium will occupy the entire classroom. This is because helium is a gas, and gases have the property of expanding to fill the entire volume of their container. In this case, the classroom acts as the container. The helium gas molecules are in constant motion and will spread out in all directions until they are evenly distributed throughout the classroom. Question 1: Which of the following best describes the behavior of gas particles? a) They are closely packed and have a fixed shape. b) They are far apart and move freely in all directions. c) They are closely packed but can flow past each other. d) They vibrate in fixed positions. Correct Answer: b) They are far apart and move freely in all directions. Question 2: If you release a gas from a small container into a larger room, the gas will: a) Stay in one corner of the room. b) Condense into a liquid. c) Spread out and occupy the entire room. d) Form a solid. Correct Answer: c) Spread out and occupy the entire room. Question 3: Why does one mole of a gas occupy a much larger volume than one mole of a liquid? a) Gas particles are larger than liquid particles. b) Gas particles are more closely packed than liquid particles. c) Gas particles are more spread out and have more space between them compared to liquid particles. d) Gases are always at a higher temperature than liquids. Correct Answer: c) Gas particles are more spread out and have more space between them compared to liquid particles. Question 4: Imagine you have two balloons of equal size, one filled with water and the other with air. Which statement is true regarding the molar volume of the substances in the balloons? a) The balloon with water has a larger molar volume. b) The balloon with air has a larger molar volume. c) Both balloons have the same molar volume. d) The molar volume depends on the temperature. Correct Answer: b) The balloon with air has a larger molar volume. Question 5: If you release helium gas into a classroom, what will happen to the distribution of helium molecules? a) They will stay close to where they were released. b) They will spread out evenly throughout the classroom. c) They will form a liquid on the floor. d) They will exit the classroom immediately. Correct Answer: b) They will spread out evenly throughout the classroom. Concept 3: Molar Mass and Molar Volume As the molar mass of the gas increases, the molar volume of a gas at Standard Temperature and Pressure (STP) decreases. 🔎 Example: Oxygen gas (O2) has a higher molar mass than helium gas (He). Therefore, one mole of oxygen gas occupies a smaller volume than one mole of helium gas at STP. 📝 Quick Question: Which gas occupies a smaller volume at STP, nitrogen (N2) or methane (CH4)? 🌟 Answer : As the molar mass of a gas increases, the individual gas particles have more mass. However, one mole of any gas at STP occupies the same volume (22.4 liters). This means that in a given volume, there is more mass packed into the same space, effectively decreasing the molar volume. 3 MCQs to test understanding for Concept 3: Question 1 (Concept 3): Which gas has a smaller molar volume at STP? a) Helium (He) b) Oxygen (O2) c) Both have the same molar volume d) Cannot be determined Correct Answer: b) Oxygen (O2) Question 2 (Concept 3): If Gas A has a higher molar mass than Gas B, which statement is true at STP? a) Gas A has a larger molar volume than Gas B. b) Gas A has a smaller molar volume than Gas B. c) Both gases have the same molar volume. d) The molar volume is independent of molar mass. Correct Answer: c) Both gases have the same molar volume. Question 3 (Concept 3): At STP, one mole of neon gas (Ne) occupies: a) A larger volume than one mole of argon gas (Ar). b) A smaller volume than one mole of argon gas (Ar). c) The same volume as one mole of argon gas (Ar). d) Twice the volume of one mole of argon gas (Ar). Correct Answer: c) The same volume as one mole of argon gas (Ar). Now, let's create 2 MCQs that combine Concepts 1, 2, and 3: Question 4 (Combining Concepts 1, 2, and 3): If a gas is released from a balloon into a room at STP, which of the following is true? a) The gas will occupy a smaller volume in the room than it did in the balloon. b) The gas will occupy a larger volume in the room than it did in the balloon. c) The molar volume of the gas will decrease. d) The gas will condense into a liquid. Correct Answer: b) The gas will occupy a larger volume in the room than it did in the balloon. Question 5 (Combining Concepts 1, 2, and 3): At STP, which of the following has the largest molar volume? a) One mole of a gas b) One mole of a liquid c) One mole of a solid d) All have the same molar volume Correct Answer: a) One mole of a gas Concept 4: Atomicity and Molar Volume As the number of atoms per molecule (atomicity) increases, the molar volume of a gas at STP decreases. 🔎 Example: Oxygen gas (O2) has two atoms per molecule, while ozone (O3) has three. One mole of ozone occupies a smaller volume than one mole of oxygen gas at STP. 📝 Quick Question: Which gas occupies a larger volume at STP, neon (Ne) or argon (Ar)? 🌟 Answer : As the number of atoms per molecule increases, the complexity and size of the molecule also increase. However, one mole of any gas at STP occupies the same volume (22.4 liters). This means that in a given volume, there are more atoms packed into the same space when the atomicity is higher, effectively decreasing the molar volume. Comparison of Molar Volumes This picture will visually represent how the molar volume changes with molar mass and atomicity. 📘 Question 1 (Combining Concepts 1, 2, 3, and 4) 📘 🔹 At STP, which gas has a smaller molar volume? a) Oxygen (O2) b) Ozone (O3) c) Both have the same molar volume d) Cannot be determined 🌟 Correct Answer: b) Ozone (O3) 🌟 📘 Question 2 (Combining Concepts 1, 2, 3, and 4) 📘 🔹 If you release a gas with high atomicity into a room, what will happen to the molar volume of the gas compared to when it was in a container? a) The molar volume will increase. b) The molar volume will decrease. c) The molar volume will remain the same. d) The molar volume will become zero. 🌟 Correct Answer: a) The molar volume will increase. 🌟 📘 Question 3 (Combining Concepts 1, 2, 3, and 4) 📘 🔹 Which of the following gases has the largest molar volume at STP? a) Helium (He) b) Methane (CH4) c) Both have the same molar volume d) Cannot be determined 🌟 Correct Answer: c) Both have the same molar volume. 🌟 📘 Question 4 (Combining Concepts 1, 2, 3, and 4) 📘 🔹 What happens to the molar volume of a gas as the number of atoms in its molecules increases, while keeping the temperature and pressure constant? a) The molar volume increases. b) The molar volume decreases. c) The molar volume remains the same. d) The molar volume becomes zero. 🌟 Correct Answer: b) The molar volume decreases. 🌟 📘 Question 5 (Combining Concepts 1, 2, 3, and 4) 📘 🔹 If a gas is compressed into a smaller container at constant temperature, what happens to its molar volume? a) It increases. b) It decreases. c) It remains the same. d) It becomes zero. 🌟 Correct Answer: b) It decreases. 🌟 Summary In this lesson, we learned about the molar volume of gases and how it is influenced by the molar mass and atomicity. We also understood that gases occupy the volume of their container. Understanding these concepts is fundamental in studying the behavior of gases in chemistry. Quiz Time! 📘 Part 1: Molar Volume of Gases Quiz 📘 🔹 Question 1 🔹 [10 Marks] 🌟 Which of the following statements accurately describes the behavior of gas particles? a) They are closely packed and have a fixed shape. b) They are far apart and move freely in all directions. c) They are closely packed but can flow past each other. d) They vibrate in fixed positions. 🔹 Question 2 🔹 [10 Marks] 🌟 What is the molar volume of any gas at Standard Temperature and Pressure (STP)? a) 1 liter 22.4 liters 6.022 x 10²³ liters 1000 liters 🔹 Question 3 🔹 [10 Marks] 🌟 True or False: As the molar mass of a gas increases, the molar volume of the gas at STP also increases. 🔹 Question 4 🔹 [10 Marks] 🌟 Which gas has a smaller molar volume at STP? a) Oxygen (O2) b) Helium (He) c) Both have the same molar volume d) Cannot be determined 🔹 Question 5 🔹 [10 Marks] 🌟 If you release a gas with high atomicity into a room, what will happen to its molar volume compared to when it was in a container? 🔹 Question 6 🔹 [10 Marks] 🌟 Which of the following gases has the largest molar volume at STP? a) Carbon dioxide (CO2) b) Nitrogen (N2) c) Methane (CH4) d) All have the same molar volume 🔹 Question 7 🔹 [10 Marks] 🌟 True or False: The molar volume of a gas at STP depends on its temperature. 🔹 Question 8 🔹 [10 Marks] 🌟 What happens to the molar volume of a gas as the number of atoms in its molecules increases, while keeping the temperature and pressure constant? 🔹 Question 9 🔹 [10 Marks] 🌟 At STP, one mole of helium gas (He) occupies: a) A larger volume than one mole of argon gas (Ar). b) A smaller volume than one mole of argon gas (Ar). c) The same volume as one mole of argon gas (Ar). d) Twice the volume of one mole of argon gas (Ar). 🔹 Question 10 🔹 [10 Marks] 🌟 If a gas is compressed into a smaller container at constant temperature, what happens to its molar volume? Thank you for your attention and participation in today's lesson! If you have any questions or need further clarification on any of the concepts, please don't hesitate to ask. Happy learning! 📚🔬 Quiz Answer Here Concepts 1 to 4 Final Quiz - Answers 📘 🔹 Question 1 🔹 [10 Marks] 🌟 Correct Answer: b) They are far apart and move freely in all directions. 🔹 Question 2 🔹 [10 Marks] 🌟 Correct Answer: b) 22.4 liters 🔹 Question 3 🔹 [10 Marks] 🌟 Correct Answer: False 🔹 Question 4 🔹 [10 Marks] 🌟 Correct Answer: c) Both have the same molar volume 🔹 Question 5 🔹 [10 Marks] 🌟 Correct Answer: The molar volume will increase. 🔹 Question 6 🔹 [10 Marks] 🌟 Correct Answer: d) All have the same molar volume 🔹 Question 7 🔹 [10 Marks] 🌟 Correct Answer: True 🔹 Question 8 🔹 [10 Marks] 🌟 Correct Answer: The molar volume decreases. 🔹 Question 9 🔹 [10 Marks] 🌟 Correct Answer: c) The same volume as one mole of argon gas (Ar). 🔹 Question 10 🔹 [10 Marks] 🌟 Correct Answer: It decreases. Summary of all chapter Concepts The molar volume in the gaseous state is much larger (about 1000 times larger) than the liquid state. The volume of a gas is the volume of the container it occupies. As the molar mass of the gas increases the molar volume of a gas at STP decreases. As the number of atoms per molecule (atomicirt) increases the molar volume of a gas at STP decreases.

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⚛️ Lesson 5 ⚛️ < Back Atomic Structure Lesson 5 ⚛️ Lesson 5 ⚛️ Discover the secrets of isotopes in this visually enhanced content. Learn about their similarities and differences, how to identify them, and their impact on chemical and physical properties. Build on your understanding of atomic structure to explore the intriguing world of isotopes and unlock new dimensions of exploration and discovery. Previous Next ⚛️1.1.5 Isotopes⚛️ ✨🔬 Unveiling the Secrets of Isotopes: Similar Yet Different 🔬✨ 🌟 The Isotope Dance: Same Protons, Different Neutrons 🌟 Isotopes are like siblings within the atomic family—they share the same number of protons and electrons but have a unique twist: a different number of neutrons. 🧑🔬⚛️ To identify an isotope, we use the chemical symbol (or word) of the element, followed by a dash and the mass number. For example, carbon-12 and carbon-14 are isotopes of carbon with 6 and 8 neutrons, respectively. 🎭 💥 Chemical Properties: A Common Chemistry 💥 When it comes to chemical properties, isotopes of the same element exhibit strikingly similar behaviors. Why? It's all about the electrons! The number of electrons in their outer shells determines an atom's chemistry, and isotopes share the same number of electrons in their respective elements. 🌌🔍 Whether it's carbon-12 or carbon-14, their outer electron shells hold the same number of electrons. Thus, they participate in chemical reactions in the same way, showcasing identical chemical characteristics. 🌟⚗️ 🌈 Physical Properties: Nuanced Differences 🌈 While isotopes share similar chemical behavior, their physical properties present subtle distinctions. The key variance lies in the number of neutrons. Neutrons are neutral subatomic particles that contribute to an atom's mass without affecting its charge. 💪 Due to these additional neutrons, isotopes exhibit slight differences in physical properties such as mass and density. These disparities, though small, are the fingerprints that set isotopes apart, enabling us to distinguish them and study their unique characteristics. ✋📊 🧠 Prerequisite: Atomic Structure 🧠 To grasp the concept of isotopes fully, understanding the fundamentals of atomic structure is crucial. This includes knowledge of protons, neutrons, and electrons, their charges, and their roles within the atom. With this foundation, we can explore the fascinating world of isotopes and their properties. 🌌💡 So, as we unveil the secrets of isotopes, remember that while they may appear similar in the world of chemistry, their underlying differences open up a whole new dimension of exploration and discovery! 🌟🚀

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Lesson 41 More Families of Elements & Periodic Trends Previous All Content Next Chapter 7 SABIS Grade 10 Part 3 Lesson 41 More Families of Elements & Periodic Trends Points explained Meaning of the word “stable” Reactions of the alkali metals with chlorine Reactions of the alkali metals with water Reactions of the alkali metals with hydrogen Flame test for Li+ , Na+ and K+ Summary of chemistry of the alkali metals 7.6 The Halogens 7.6.1 Physical properties of the halogens 7.6.2 Covalent bonding in the halogens Defining a covalent bond Differences between covalent and ionic bonding 7.6.3 Boiling points and melting points of the halogens 7.6.4 Atomic radii and volumes 7.6.5 Chemistry of the halogens Reactions with the alkali metals Summary 7.6.6 Chemistry of the halide ions The halides are stable Test for the halide ions Relative reactivity of the halogens 7.7 Hydrogen - A Family by Itself 7.7.1 Physical properties 7.7.2 Chemistry of hydrogen Reaction with the alkali metals 7.8 The Third-row Elements 7.8.1 Physical properties of the third-row elements 7.8.2 Compounds of the third-row elements The hydrides The chlorides The oxides Summary 7.9 The periodic table: chemical reactivity 🔬Understanding Stability, Alkali Metals & Halogens 📚Pre-Requisite Questions: What does it mean for an element to be "stable"? 🤔 What happens when alkali metals react with chlorine? 🧪 Can you describe the flame test results for Li+, Na+, and K+? 🔥 Break for Reflection 🤔✍️ (Answers: 1. A stable element has a full outer electron shell and doesn't tend to react. 2. When alkali metals react with chlorine, they form ionic salts. 3. Li+ burns with a crimson flame, Na+ with a yellow flame, and K+ with a lilac flame.) 🚀 Lesson Begins! 🧱 Meaning of the Word “Stable” In the chemistry world, "stable" doesn't mean standing still! It means an atom has a full outer shell of electrons and is not looking to react. They're like that chill friend who's content with what they have! 😌 💥Reactions of the Alkali Metals 💦With Water Splash alert! Alkali metals react violently with water, producing heat, hydrogen gas, and an alkali metal hydroxide. Think of it as a bath bomb that's too explosive for the tub! 🛀💣 🎈With Hydrogen Pairing up! Alkali metals can combine with hydrogen to form metal hydrides, releasing energy in the process. It's like an energetic dance duo! 💃🕺 🔥Flame Test for Li+, Na+, and K+ Ready for some fireworks? 🎆 In the flame test, Li+ produces a red/crimson flame, Na+ gives a yellow flame, and K+ presents a lilac flame. It's like a mini festival of lights in the lab! 🎇 7.6 The Halogens The Halogens, just like the Alkali Metals, are an interesting bunch. They're like the goths of the periodic table, always looking to gain an electron to achieve stability. 🕶️💀 🧪Covalent Bonding in the Halogens Halogens form covalent bonds by sharing electrons. Imagine sharing your favorite pizza with a friend—that's how halogens share electrons to become stable. 🍕❤️ 🔥Boiling Points and Melting Points of the Halogens Halogens have higher boiling and melting points as we move down the group, thanks to the increasing number of electrons which cause stronger intermolecular forces. It's like adding more logs to the fire—the more you have, the higher the flame! 🏕️🔥 ⚖️Atomic Radii and Volumes Atomic radii also increase as we go down the group. It's kind of like siblings—the older ones tend to be bigger! 🧑🤝🧑 🔥Chemistry of the Halogens Halogens are pretty reactive. Their reactions with alkali metals form ionic salts, and they're not shy about displacing less reactive halogens. It's like a game of musical chairs! 🎶🪑 7.7 Hydrogen - A Family by Itself Hydrogen is unique. Despite being the lightest and simplest element, its properties don't quite fit into any group. So, it charts its own path—just like a lone wolf. 🐺⛰️ 7.8 The Third-row Elements The third-row elements are like the middle kids of the periodic table. They have their quirks and surprises! So, let's dive deeper into their physical properties and compounds. 🏊♂️🌊 7.9 The Periodic Table: Chemical Reactivity The periodic table is not just a chart; it's a tale of reactivity, trends, and atomic friendships. Keep exploring, keep learning! 🚀 Review Questions: What is meant by a stable element? a. It has a full outer electron shell b. It has no protons c. It is radioactive d. None of the above What happens when alkali metals react with water? a. Nothing b. They dissolve c. They produce heat and hydrogen gas d. They turn into halogens Which element doesn't fit well into any group in the periodic table? a. Hydrogen b. Helium c. Oxygen d. Nitrogen Which of the following is NOT a property of the halogens? a. They form ionic bonds b. They form covalent bonds c. They have high boiling and melting points d. They are very reactive (Answers: 1. a, 2. c, 3. a, 4. a) End of Lesson 3 ⭐Keep studying, keep learning!⭐

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