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Percentage Yield | Chemical Calculations | Chemistry |  FuseSchool

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How To Calculate Relative Atomic Mass | Chemical Calculations | Chemistry | FuseSchool

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• RELATIVE MASSES
• As covered in Topic 1, the relative atomic mass (Ar) is the mass of an atom compared to that of carbon-12
• Calculating the relative formula mass (Mr):
• The ‘relative formula mass’ (Mr) of a substance is the sum of the relative atomic masses (Ar) of all the atoms or ions in its formula:
• E.g CO2 – 1 carbon atom, 2 oxygen atoms:
o Ar of C = 12
o Ar of O = 16… x 2 = 32
o relative formula mass (Mr) of CO2 = 12 + 32 = 44
• E.g2 Ca(NO3)2 – 1 calcium atom, 2 nitrogen atoms, 6 oxygen atoms:
o Ar of Ca = 40
o Ar of N = 14… x 2 = 28
o Ar of O = 16… x 6 = 96
o relative formula mass (Mr) of Ca(NO3)2  = 40 + 28 + 96 = 164
• Molecular and empirical formulae:
• The true formula for a simple molecular compound is called the ‘molecular formula’ – this shows the actual number of atoms of each element in a molecule
• Substances can also be represented by an empirical formula – this shows the simplest whole number ratio of atoms or ions of each element in a substance
• E.g:
o Ethene has the molecular formula C2H4
o Propene has the molecular formula C3H6
o For every molecule of ethene or propene, there are twice as many hydrogen atoms than carbon atoms…
 the molecular formulae of both ethene and propene can be simplified to CH2 – the empirical formula
o So ethene and propene have different molecular formulae (C2H4 and C3H6) but the same empirical formula (CH2)
• Calculating the empirical formula:
• If you know the mass (in grams) of each element present in a compound, you can use this and the relative atomic mass of each element to calculate the empirical formula of the compound……e.g for calcium chloride:
Ca         Cl
Mass in g                   10.0       17.8
Relative atomic mass      40       35.5
Step 1. Divide the mass of each element by its relative atomic mass 10.0/40 = 0.25 17.8/35.5 = 0.5
Step 2. Divide the answers by the smallest number obtained after step 1 (in this case 0.25) to find the simplest ratio 0.25/0.25 = 1 0.5/0.25 = 2
i.e for every molecule of calcium chloride, there are twice as many chlorine atoms than calcium atoms
Empirical formula CaCl2
• Note: you can just use charges of ions to work out the molecular formula (see Topic 2) and then simplify it as far as possible to get the empirical formula
• PERCENTAGE COMPOSITION
• You can use the relative masses to calculate the percentage composition of each element in a compound…
• Percentage (by mass) of an element in a compound = number of atoms of element x Ar/Mr x 100…
• E.g calculate the percentage by mass of oxygen in potassium nitrate – KNO3…
o Number of atoms of oxygen in KNO3 = 3
o Relative atomic mass (Ar) of oxygen = 16
o Then calculate the relative formula mass (Mr) of KNO3…
 KNO3 – 1 potassium atom, 1 nitrogen atom, 3 oxygen atoms:
 Ar of K = 39
 Ar of N = 14
 Ar of O = 16…x 3 = 48
 Mr of KNO3 = 101
o % of oxygen in KNO3 = 3 x (16/101) x 100 = 47.5%
• Calculating the masses of reactants or products:
• During a chemical reaction, no atoms are lost or made – they are just rearranged to make new substances (the ‘products’)
• You can use relative masses and the balanced equation for a reaction to calculate the mass of a reactant or product…
• E.g potassium nitrate (KNO3) is decomposed to potassium nitrite (KNO2) and oxygen (O2)......what mass of potassium nitrate is needed to make 1.6g of oxygen?
o Step 1 – balanced chemical equation: 2KNO3  2KNO2 + O2
o Step 2 – work out the relative masses of the substances needed in the calculation:
 Mr of KNO3 (see above) = 101
 There are two particles of KNO3 in the balanced equation we multiply the Mr of KNO3 by 2… 101 x 2 = 202
 Relative mass (only one atom so Ar/Mr it’s the same) of O2 = 32
o Step 3 – divide the answers by the smallest relative mass calculated above (in this case 32) to find the ratio:
 2KNO3: 202/32 = 6.3
 O2: 32/32 = 1
 ratio of 2KNO3:O2 is 6.3:1
o Step 4 – use the ratio to find the answer!
 Ratio is 6.3:16.3g of potassium nitrate are needed to make 1g of oxygen
 to find the mass of potassium nitrate needed to make 1.6g of oxygen we multiply 1.6 by 6.3……1.6 x 6.3 = 10.1
 10.1g of potassium nitrate are needed to make 1.6g of oxygen
• YIELDS
• The amount of useful product that is obtained from a chemical reaction is called the yield
• In theory one might expect all the reactants to turn into products – this is the ‘theoretical yield’
• The theoretical yield can be calculated from the balanced equation of a reaction…
• Calculating the theoretical yield:
• E.g 2H2 + O2  2H2O…:
o Calculate the relative formula masses:
 2H2 – 4 hydrogen atoms:
 Ar of H = 1……Ar of 2H2 = 2 x 2 = 4g
 O2 – 2 oxygen atoms:
 Ar of O = 16……Ar of O2 = 32g
 2H2O – 4 hydrogen atoms, 2 oxygen atoms
 32 + 4 = 36g
o So 36g of water should theoretically be produced when 4g of hydrogen reacts with 32g of oxygen
• Percentage yield:
• In practice, however, the ‘actual yield’ obtained is less than the predicted ‘theoretical yield’ (see reasons for this below)
• The percentage yield compares the actual yield to the theoretical yield (i.e it compares the actual amount of product formed to the predicted amount of product formed calculated from the balanced chemical equation)…:
o Percentage yield = (actual yield / theoretical yield) x 100
o E.g if in the reaction above the actual yield (i.e the amount of water produced) was in fact 30g…percentage yield = (30/36) x 100 = 83.3%
• Why is the yield less than expected?
• The theoretical yield of a reaction assumes that…:
o all the reactants are turned into products
o the products are successfully separated from the reaction mixture
• 3 main reasons for why reactions don’t give 100% (percentage) yields:
o 1. reaction may be incomplete – i.e not all reactants are used up
o 2. some of the product is lost during the practical preparation – e.g when transferring liquids from one container to another
o 3. there may be other unwanted reactions taking place – e.g some of the reactants may react in different ways to make a different product
• WASTE AND PROFIT
• Disposal of waste products:
• Many chemical reactions produce substances other than the substance that is wanted - these additional substances are called by-products
• Some by-products can be useful…:
o E.g sodium hydroxide and hydrogen – by-products of the electrolysis of sodium chloride solution – are useful because they can be sold
• However, many by-products are useless – useless by-products are called waste products and have to be got rid of…
• Disposal of waste products can…:
o be expensive – the waste may have to be transported to a landfill site or it may have to be treated with another substance to make it safe
o cause environmental problems
o cause social problems…e.g:
 house prices could drop if a chemical plant is built near to them
 unpleasant smells can be emitted from landfill sites
• Finding the most cost effective process:
• As well as taking the environmental impact into account, chemical companies want to use reactions that will make the most money (profit)
• To do this they try to use reactions in which…:
o the percentage yield is high – this reduces costs
o the reaction takes place quickly – this reduces costs
o all the products of the reaction are usefulno waste products
• E.g when ammonia is made there are no waste products:
o N2 + 3H2  2NH3
• E.g2 when iron is extracted from iron oxide there is always waste carbon dioxide:
o Fe2O3 + 3CO  2Fe + 3CO2
o However, if use can be found for waste products then a reaction becomes commercially viable (i.e money/profit can be made from it)

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GCSE BBC Bitesize Higher Science - Exothermic and Endothermic Chemical reactions

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This Higher GCSE BBC Bitesize video is from the original programmes from 2000 that were broadcast on BBC2. It covers the areas of the Chemistry Higher paper. Select the, "more from," or type jamjarmmx into your search for the other Higher Chemistry clips as well as the Physics and Biology clips.

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Rates of Reaction - Part 2 | Reactions | Chemistry | FuseSchool

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How to speed up chemical reactions (and get a date) - Aaron Sams

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• TEMPERATURE CHANGES
• In chemical reactions, transfer of heat energy between the reactants and the surroundingschange in temperature
• Exothermic reactions:
• Exothermic reactions give out heat energy to the surroundingstemperature of the reaction mixture and its surroundings increases
• Most reactions are exothermic…
• Examples of exothermic reactions:
o All combustion reactions e.g combustion of methane:
 methane + oxygen  carbon dioxide + water
o Explosions (release a lot of heat energy very quickly)
• Endothermic reactions:
• Endothermic reactions take in heat energy from the surroundingstemperature of the reaction mixture and of the surroundings decreases
• Few chemical reactions are endothermic…
• Examples of endothermic reactions:
o Reaction of sodium hydrogencarbonate with hydrochloric acid (if you touch the test tube it feels cold because heat energy is taken from your hand)
o Dissolving ammonium nitrate in water
o Photosynthesis (takes in heat energy from the sun)
• Making and breaking bonds:
• In chemical reactions, bonds in the reactants are broken and new bonds are formed to make the products:
o Breaking bonds requires energyit’s an endothermic process
o Making bonds releases energyit’s an exothermic process
• In an exothermic reaction:
o More heat energy is released making bonds (in the products) than is required to break bonds (in the reactants)
o heat energy is released
• In an endothermic reaction:
o More heat energy is required to break bonds (in the reactants) than is released making bonds (in the products)
o heat energy is absorbed
• These energy changes can be shown using diagrams:

• RATES OF REACTION
• The rate of a chemical reaction is the speed at which it takes place – i.e it tells us how fast reactants are used up and products are made:
o E.g explosions are very fast chemical (exothermic) reactions that release lots of heat energy and gas in a short space of time
o E.g2 rusting of iron is an everyday example of a slow chemical reaction
• Collision theory:
• For a reaction to occur, particles of the reactants must collide with each other
o if particles collide more frequently, rate of reaction will increase
• However, not all collisions lead to a reaction – particles must collide with enough energy to break the bonds in the reactants:
o the more energy particles have, the greater the chance of ‘successful collisions’ (i.e greater chance of particles reacting when they collide)
o rate of reaction will increase if there are more high-energy collisions between particles
• The effect of temperature:
• The higher the temperature of the reactants, the faster the reaction - e.g eggs cook faster in boiling water than in warm water
• Explanation:
o As temperature is increased, particles move faster…:
 1. collisions are more frequent
 2. particles have more energygreater chance of successful collisions
 increased rate of reaction
• Sometimes we cool reactions to slow them down:
o E.g some foods are put in a fridge to slow down chemical reactions that make food go off
• The effect of concentration:
• The higher the concentration of a reactant in a solution, the faster the reaction
• Explanation:
o The more concentrated a solution, the more solute particles there are in a given volume:
 more likely that reactant particles will collide with one another
 increased rate of reaction
• The effect of surface area:
• The greater the surface area of a solid, the faster the rate of reaction
• Explanation:
o Increasing the surface area increases the number of particles exposed on the surface (that can collide and react)
o more frequent collisions between reactant particles
o increased rate of reaction
• To increase surface area, solids are crushed into lots of small pieces:
o E.g in power stations, coal is ground up into a fine powder to help it burn faster
• CATALYSTS
• Catalysts are substances that speed up chemical reactions without being used up in the reactions (note: they are neither reactants nor products)
• Catalysts speed up chemical reactions by increasing the probability of successful collisions - i.e when particles collide they don’t need as much energy to react
• Many chemical processes use catalysts to increase the rate at which products are generated
• Using catalysts means that reactions can be done at lower temperatures and pressures than they would otherwiseless energy is usedsaves money
• Catalytic converters:
• The combustion of petrol in car engines produces carbon monoxide (toxic) and unburned hydrocarbons
• cars are now built with catalytic converters:
o Catalytic converters help to combine carbon monoxide and unburned petrol with oxygen from the air to form carbon dioxide and water vapourreducing pollutants in exhaust gases
o Catalytic converters contain the transition metals platinum, rhodium or palladium (all very expensive), which act as catalysts
o The catalysts have a high surface area (‘honeycomb structure’)…:
 1. rate at which carbon monoxide and unburned petrol react with oxygen from the air to form carbon dioxide and water vapour is increased
 2. less of the expensive metals are usedcosts are lower
o Catalytic converters work best at high temperatures (as particles collide more frequently and with more energy)…:
 When a car engine is first started, a catalytic converter is cool and doesn’t work well
 However, the hot gases from the engine quickly heat it up

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Group 7 - The Halogens | Properties of Matter | Chemistry | FuseSchool

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Group 1 - The Alkali Metals | The Periodic Table | Properties of Matter | Chemistry | FuseSchool

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What Are Metallic Bonds | Properties of Matter | Chemistry | FuseSchool

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Atomic Hook-Ups - Types of Chemical Bonds: Crash Course Chemistry #22

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• Properties of metallic substances:
• The atoms in metals are held together by metallic bonding (see below) – this gives metals different properties from other types of substances…:
o Metals are good conductors of electricity and heat
o Metals are solids at room temperature (metallic bonds are strong), except for mercury, which is a liquid at room temperature
o Metals don’t dissolve in water
o Metals are malleable (i.e can be hammered into shape)
• metals have many uses…e.g they’re used to make cars, buildings, tools etc…
• METALLIC BONDING AND TRANSITION METALS
• Metallic bonding:
• Metal atoms form positive ions, which are held close together in a regular arrangement by a ‘sea’ of outer shell electrons:
o The term ‘sea of electrons’ is used because electrons in the outer shells of metal atoms are free to move through the structure
o The electrons aren’t located in specific atomsthey’re called ‘delocalised electrons’

• Metals conduct electricity:
o Delocalised electrons move around randomly in all directions between the positive ions
o If a potential difference (i.e voltage) is applied across a piece of metal, all the delocalised electrons start to move in the same direction
o This movement of electrons is an electric current
• Metals are malleable:
o If a large force is applied, the layers of positive ions in a metal can slide over each other
o The positive ions are still held together by the sea of electronsthe metal spreads out (changes shape) instead of breaking
• Transition metals:
• Most metals are transition metals
• Most transition metals have high melting and boiling points and form coloured compounds
• Transition metals are in the central block of the periodic table…:

• ALKALI METALS
• The alkali metals are found in group 1 of the periodic table:
o they have 1 electron in their outer shell…
o to end up with a full outer shell they must lose their outer electron
o they form ions with a charge of +1
• The atoms in alkali metals are held together by metallic bonding
• Alkali metals are solids at room temperature, but they have low melting points compared to other metals
• Alkali metals are soft metalscan be cut with a knife
• All alkali metals react with water to form a metal hydroxide and hydrogen gas:
o E.g lithium + water  lithium hydroxide + hydrogen
o 2Li (s) + 2H2O (l)  2LiOH (aq) + H2 (g)
o Note: all metal hydroxides are alkaline
• Reactivity:
• If lithium is dropped in water, it just floats on the water and fizzes (that’s the hydrogen being produced) until the reaction is finished
• Sodium reacts more strongly with water:
o Sodium has a lower melting pointthe reaction produces enough heat to melt the metal
o This forms a molten ball of sodium that whizzes around the surface of the water (releasing hydrogen) until the reaction is finished
• Potassium reacts even more strongly and the hydrogen produced during the reaction catches fire, producing a lilac flame
• So the reactivity of alkali metals increases as you go down group 1
• Explanation for this:
o The elements at the bottom of group 1 have more electrons than the elements at the top of the groupthey have more electron shellsthe electron in the outer shell is further from the nucleus
o The attraction between positive and negative charges (i.e between the nucleus and the outer electron) is weaker when charges are further apart
o outer electron in a potassium atom (configuration: 2.8.8.1) is not held as strongly as the outer electron in a lithium atom (configuration: 2.1)
o Metals react by losing their outer electron, forming ions with a +1 charge and a full outer shell
o Potassium loses its outer electron more easilyit’s more reactive
• HALOGENS
• The halogens are the elements in group 7:
o have 7 electrons in their outer shell…
o to form a full outer shell they must gain one electron
o they form ions with a charge of -1
• At room temperature…:
o Fluorine is a pale yellow gas
o Chlorine is a yellow-green gas
o Bromine is a brown liquid
o Iodine is a grey solid
• Reactivity of the halogens:
• The pattern of reactivity for the halogens is the opposite to that of the alkali metals - i.e halogens become less and less reactive as you go down group 7:
o This is because halogens react by gaining an electron - this is easier with fewer electron shells (because outer electrons are closer to the nucleus)
o So fluorine is the most reactive halogen
• HALOGEN REACTIONS
• Reactions with metals:
• All halogens react with metals to form metal halides (the word ‘halide’ means that the compound contains only metal ions and ions of one of the halogens):
o E.g potassium + bromine  potassium bromide
o 2K (s) + Br2 (l)  2KBr (s)
• Note: chlorine forms ‘chloride’, fluorine forms ‘fluoride’, iodine forms ‘iodide’
• Reactions with hydrogen:
• Halogens react with hydrogen gas to form hydrogen halides:
o E.g hydrogen + chlorine  hydrogen chloride
o H2 (g) + Cl2 (g)  2HCl(g)
• Note: hydrogen halides form acids when they dissolve in water (e.g when hydrogen chloride dissolves in water it forms hydrochloric acid)
• Displacement reactions:
• More reactive halogens can ‘displace’ less reactive halogens from their compounds
• Less reactive halogens cannot ‘displace’ more reactive halogens from their compounds (in this case, no reaction takes place)
• displacement reactions can be used to work out the reactivity of different halogen elements…
• E.g chlorine is more reactive than bromine (as it is higher up group 7)will displace bromine from a bromide…:
o sodium bromide + chlorine  sodium chloride + bromine
o 2NaBr (aq) + Cl2 (g)  NaCl (aq) + Br2 (aq)
o Note: the solution is initially colourless but will turn an orange-brown colour as bromine is formed
• E.g2 if chlorine is added to a compound of fluorine there will be no reaction because chlorine is less reactive than fluorine (can’t displace it)
• NOBLE GASES
• The noble gases are in group 0 of the periodic table
• Elements in other groups gain full outer shells by forming ions (i.e by losing or gaining electrons) or by forming covalent bonds (i.e by sharing electrons)
• However, noble gases already have full outer shellscompared to other elements, noble gases are inert (very unreactive)
• Properties of noble gases:
• All noble gases are gases at room temperaturethey have low boiling points
• Particles in a gas are spread far apartnoble gases have low densities
• As you down group 0:
o the boiling points of noble gases increase
o the densities of noble gases increase
• Discovery of the noble gases:
• Chemists noticed that the density of pure nitrogen made in chemical reactions was less than the density of nitrogen extracted from the air
• It was hypothesised that the nitrogen extracted from air also contained a denser gas
• Experiments were done to find the identity of this dense gas – it turned out to be argon
• The other noble gases were discovered soon after
• Uses of the noble gases:
• The noble gases are useful because they are unreactive:
o Xenon and argon were used inside filament lamps, instead of air, to stop the hot filament reacting with oxygen and burning away
o Argon and helium are used in welding - they form a blanket over the hot metalpreventing it from reacting with oxygen in the air
• Argon is non-flammableused in fire-extinguisher systems
• Helium has a low densityused for filling balloons and airships
• When electric current is passed through a tube filled with neon under low pressure, coloured light is producedneon is often used in fluorescent lamps and advertising displays

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Ionic Compounds & Their Properties | Properties of Matter | Chemistry | FuseSchool

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Giant Ionic Structures or Lattices | Properties of Matter | Chemistry | FuseSchool

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What Are Covalent Bonds | Properties of Matter | Chemistry | FuseSchool

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Properties of Diamond (With Narration) | Mocomi Kids

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http://mocomi.com presents: Properties of Diamond Find out from this animated learning module about the properties that come together to form a diamond. Similar to diamonds, emeralds and rubies are too made up of a number of properties. Can you find out what they are? Diamond is an allotrope of Carbon.

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Uploaded by LaunchAt Viral Media on 2014-09-24.

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0:50 Miscible and Immiscible substances 1:39 Separation Techniques - Separating Immiscible Liquids using a separating funnel 1:39 Separation Techniques - Distilliation 4:05 Separation Techniques - Fractional Distillation of Air 5:31 Separation Techniques - Chromatography 8:00 Interpreting Chromatograms 9:44 Calculating Rf values

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Tutorial: Paper & Thin Layer Chromatography | Chemical Tests | Chemistry | FuseSchool

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Topic 3 – Covalent compounds and separation techniques

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Topic 3 – Covalent compounds and separation techniques

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COVALENT BONDS
• Most non-metal compounds are held together by pairs of electrons, which form covalent bonds
• The electrons used in covalent bonds come from the outermost shells of atoms
• When covalent bonds form, the sharing of electrons allows both atoms to have full outer shells (more stable)formation of molecules
• Covalent bonding between atoms can be shown by dot-cross diagrams (often only outer electrons are shown because it is these that form covalent bonds)…
• Covalent bonds between atoms of the same (non-metal) element:
• E.g two hydrogen atoms (H) form one molecule of hydrogen gas (H2):
(note: compounds are molecules containing at least two different elements)
o Each hydrogen atom contributes one electron to the covalent bond (i.e electrons are shared)both hydrogen atoms have full outer shells
o stable H2 molecule is formed

• Note: in a dot-cross diagram there’s no difference between the electrons from different atoms…they’re shown as dots and crosses just to show which atom each electron is from
• Covalent bonds between atoms of different (non-metal) element
Double bonds:
• Atoms can share more than one pair of electrons if this is needed for each atom to have a full outer shell…
• Two pairs of shared electrons form a double bond…
o E.g O=O (O2)
o E.g2 O=C=O (CO2)

PROPERTIES OF COVALENT SUBSTANCES
• Simple molecular covalent substances:
• These include gases such as hydrogen, methane, oxygen, carbon dioxide, and liquids such as water
• These substances have low melting and boiling points because (although there are strong covalent bonds between atoms in each molecule) there are only weak forces between neighbouring molecules
• They are also poor conductors of electricity because they haven’t gained or lost electronsthere are no charged particles that can move around (this is essential for conduction of electricity to take place)
• Giant molecular covalent substances:
• E.g sand (made of silicon and oxygen atoms), diamond and graphite (both made of carbon atoms)
• These consist of billions of atoms all joined together by covalent bonds
• Most of these substances have high melting and boiling points because all the atoms are joined to other atoms by strong covalent bonds (lots of heat energy is needed to break these bonds)
• Properties of diamond and graphite:
• Both diamond and graphite have high melting and boiling points because of the strong covalent bonds between the carbon atoms
• However…:
o Carbon atoms in diamond form a compact structure
o Carbon atoms in graphite form layers
o graphite and diamond have different properties
• Diamond:
o is very hard because all the atoms are joined with strong covalent bonds… diamond is used to make cutting tools
o doesn’t conduct electricity because there are no free (‘delocalised’) electrons that can move around (i.e all four outer shell electrons in each carbon atom are involved in making bonds)
• Graphite:
o is very soft because although the covalent bonds within the layers are very strong, there are only weak forces between the layers…graphite is used as a lubricant
o conducts electricity because...:
 only 3 outer shell electrons in each carbon atom are involved in making bonds
 one electron from each carbon atom is free to move along the layers (electron is said to be ‘delocalised)current can flow
 graphite is used to make electrodes
• MISCIBLE OR IMMISCIBLE?
• Immiscible liquids:
• Liquids that don’t mix completely with each other (e.g oils in water) are ‘immiscible’
• Even when shaken up, immiscible liquids soon separate out again
• Immiscible liquids can be separated using a separating funnel:
o Tap of funnel is opened so the lower liquid runs out of the funnel and can be collected in a beaker
o Tap is then closed before the other (upper) liquid starts running out
o A different beaker is placed under the funnel and tap is opened to allow upper liquid to be collected
• Miscible liquids:
• If two liquids dissolve in each other, their particles mix completely to make a solution – liquids like this are ‘miscible’
• Once mixed, the only way to separate miscible liquids is by fractional distillation
• Fractional distillation can separate mixtures of miscible liquids because they have different boiling points…:
o The mixture of liquids is heated and the liquids evaporate
o The vapours condense in a fractionating column:
 The fraction with the highest boiling point condenses near the bottom of the column (where it’s hotter)
 The fraction with the lowest boiling point condenses near the top of the column (where it’s cooler)
• Fractional distillation can also be used to separate oxygen and nitrogen in air:
o Nitrogen boils at -183°C, oxygen boils at -196°C
o The air first has to be separated into a mixture of liquids:
 To remove water vapour in the air, the air is cooled so that the water can freeze and be removed
 The remaining air is cooled to -200°C (below boiling points of both nitrogen and oxygen)air is in a liquid state
o Nitrogen has a lower boiling point than oxygenwhen liquid air is warmed to -185°C, the nitrogen evaporates and rises up the column
o The oxygen stays as liquid and is piped out of the bottom of the column
CHROMATOGRAPHY
• Inks, paints and foods often contain mixtures of coloured compounds
• Some coloured compounds dissolve better in a solvent (a ‘solvent’ is what substances dissolve into…e.g water) than others
• mixtures of coloured compounds can be separated by their solubilities
• This is done by chromatography…:
o In paper chromatography, samples are placed near the bottom of a sheet of special paper (‘base line’ on diagram)
o The solvent (e.g water) soaks up the paper (solvent must be placed above the bottom edge of the paper but below where the samples are placed)
o More soluble compounds in a sample are carried up the paper faster (and further) than less soluble onesseparating them
o The paper with the separated components on it is called a chromatogram
• Rf value:
• The Rf value is the distance the compound has moved up the paper (‘b’ on diagram) divided by the distance the solvent has moved (‘a’ on diagram)
• The further up the paper the compound has moved, the greater the Rf value
• the greater the Rf value, the more soluble the compound
• Uses of chromatography:
• The Food Standards Agency uses chromatography to separate and identify food colourings – this ensures colourings used in foods and drinks are safe
• The police use chromatography to compare a suspect’s DNA sample to the DNA sample found at the crime scene
• Chromatography can also be used to analyse paints and dyes – this helps museum staff to mix exact copies of old-fashioned paints, to restore old paintings or to identify fakes

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Testing for positive ions - part 2 | Chemistry for All | The Fuse School

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Learn the basics about Testing for positive ions - part 2. Find out more in this video! Feel to free to review part 1 of this lesson: https://www.youtube.com/watch?v=H8awN19uVPA&hd=1 This Open Educational Resource is free of charge, under a Creative Commons License: Attribution-NonCommercial CC BY-NC ( View License Deed: http://creativecommons.org/licenses/by-nc/4.0/ ).

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Testing for positive ions - Part 1 | Chemistry for All | The Fuse School

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Learn the basics about Testing for positive ions in the first part of the lesson. How can Ions be detected? What methods or techniques are used? Find out more in this video! This Open Educational Resource is free of charge, under a Creative Commons License: Attribution-NonCommercial CC BY-NC ( View License Deed: http://creativecommons.org/licenses/by-nc/4.0/ ).

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How to test for negative ions | The Chemistry Journey | The Fuse School

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Find out all the basics of testing for negative ions. Make sure to watch our video on testing for positive ions, too: http://youtu.be/4TgW0hGnqL0 At Fuse School, teachers and animators come together to make fun & easy-to-understand videos in Chemistry, Biology, Physics, Maths & ICT. Our OER are available free of charge to anyone.

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Precipitation Reactions: Crash Course Chemistry #9

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Formulae Of Ionic Compounds & Their Names - Part 1 | Properties of Matter | Chemistry | FuseSchool

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What are Ionic Bonds?  | Properties of Matter | Chemistry | FuseSchool

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IONIC BONDS
• An ion is an atom or groups of atoms with a positive or negative charge
• Ionic bonds form between positively and negatively charged ions (i.e between cations and anions – see below)
• Atoms are most stable with full outer shells
• However, atoms of most elements have incomplete (i.e not full) outer shells they readily lose or gain electrons during chemical reactions to obtain full outer shells
• When electrons are gained or lost, the atoms become ions…
• Cations:
• Metal atoms readily lose their outermost electrons to form positively charged ions called cations
• For elements in groups 1 and 2, the number of outer electrons lost is the same as their group number…:
o E.g sodium is in group 1 has 1 electron in its outer shell (2.8.1). It can lose one electron to become a Na+ cation with a full outer shell (2.8)
o E.g2 magnesium is in group 2 has 2 electrons in its outer shell (2.8.2). It can lose two electrons to become a Mg2+ cation with a full outer shell (2.8)
• Anions:
• Non-metal atoms can gain electrons to form negatively charged ions called anions
• For elements in groups 6 and 7, the number of electrons they gain is 8 minus their group number…:
o E.g oxygen is in group 6 has 6 electrons in its outer shell (2.8.6). It can gain two electrons to become an O2- anion with a full outer shell (2.8.8)
o E.g2 chlorine is in group 7 has 7 electrons in its outer shell (2.8.7). It can gain one electron to become a Cl- anion with a full outer shell (2.8.8)
• Note: when non-metal atoms form ions, their name changes to –ide (e.g chlorine atoms are called chloride ions…oxygen atoms are called oxide ions)
IONIC COMPOUNDS
• Ionic compounds form when a metal reacts with a non-metal:
o Electrons lost by the metal are transferred to the non-metal
o both positive and negative ions that form end up with stable, full outer shells
o Oppositely charged ions attract each other stronglyforming an ionic compound held together by ionic bonds
• E.g when sodium and chlorine react…:
o Electrons are transferred from sodium atoms to chlorine atoms
o This forms Na+ and Cl- ions (both have a stable, full outer shell)
o These oppositely charged ions attract each other stronglyforming the ionic compound NaCl (table salt)
• Working out a formula:
• Ionic compounds are electrically neutral because they contain equal numbers of positive and negative charges
• If you know the charges of the cations and anions you can work out the formula of the ionic compound…
• E.g for sodium chloride, one Na+ is needed for every Cl- ionformula is NaCl
• E.g2 the ionic compound aluminium oxide:
o Aluminium oxide contains Al3+ cations and O2- anions
o For charges to be equal you must have…:
Two Al3+ cations ( + 6 charge)
Three O2- anions ( - 6 charge)
o formula of aluminium oxide is Al2O3 (neutral charge)
• Compound ions:
• Compound ions contain more than one element
• E.g the nitrate ion NO3- contains one nitrogen atom joined to three oxygen atoms plus an extra electron (negative charge)
• If two or more compound ions of the same type are needed in a formula, the ion must be written inside brackets (with the number on the outside)…
• E.g the ionic compound magnesium nitrate:
o Magnesium nitrate contains Mg2+ and NO3- ions
o For charges to be equal you need two NO3- ions for every Mg2+ ion
o formula of magnesium nitrate is Mg(NO3)2
Names of ionic compounds:
Note:
o Ionic compounds usually end in –ide (e.g NaCl - sodium chloride)
o However, compounds that contain oxygen atoms end in –ate (e.g Mg(NO3)2 - magnesium nitrate)
• The structure of ionic compounds:
• The ions in an ionic compound are packed tightly together and arranged in a regular way, called a lattice structure
• The lattice structure is held together by strong electrostatic forces of attraction (i.e by ionic bonds) between oppositely charged ions
• PROPERTIES OF IONIC COMPOUNDS
• Conducting electricity:
• Ionic compounds don’t conduct electricity when solid
• However, they do conduct electricity…:
o when molten (i.e when they’re heated until they turn into liquid)
o when in aqueous solution (i.e dissolved in water) – this is why sea water (NaCl in aqueous solution) conducts electricity
• Explanation:
o Two conditions must be met for a substance to conduct electricity…
Must contain charged particles
Charged particles must be free to move
o Ionic compounds contain charged particles (ions) but these are only free to move when molten or in aqueous solution
o In the solid form, ions can’t move (they’re strongly held together in a lattice structure by ionic bonds)can’t conduct electricity
• Melting points and boiling points:
• The melting point of a substance is the temperature at which it changes from a solid to a liquid
• The boiling point of a substance is the temperature at which it changes from a liquid to a gas (at its fastest possible rate)
• Ionic bonds holding ionic compounds together in a lattice are very strong:
o This means lots of (heat) energy is needed to break the ionic bonds
o ionic compounds have high melting and boiling points
o ionic compounds are usually solids at room temperature
SOLUBILITY
• If a substance dissolves well in a particular liquid, it is said to be ‘soluble’
• If a substance doesn’t dissolve at all in a particular liquid, it is said to be ‘insoluble’
• A ‘salt’ (not NaCl which is ‘table salt’) is a substance that can be made by reacting an acid and an alkali
• Solubility rules for salts in water:
Soluble in water Insoluble in water
All common sodium, potassium and ammonium salts
All nitrates
Most chlorides Silver/lead chlorides
Most sulfates lead/barium/calcium sulfates
sodium/potassium/ammonium carbonates Most carbonates
sodium/potassium/ammonium hydroxides Most hydroxides
• Precipitation reactions:
• A reaction in which at least one insoluble solid (called the ‘precipitate’) is produced from two soluble substances is called a precipitation reaction:
o E.g lead nitrate + potassium iodide --> lead iodide + potassium nitrate
o Balanced equation: Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)
o State symbols show that all substances are dissolved in water (aq - aqueous) except for lead iodide, which is insoluble (so it’s shown as a solid – s) forms a precipitate
• In precipitation reactions, the precipitate can be separated from the unreacted ions by filtration. It is washed on filter paper and then dried in a warm oven
PRECIPITATES
• Using the solubility rules, it’s possible to work out which precipitate will form when two solutions are mixed together
• E.g what precipitate is formed when copper chloride and potassium hydroxide are mixed together?
o When salts mix, the ions swap so that: copper chloride + potassium hydroxide  copper hydroxide + potassium chloride
o All potassium salts are soluble, whereas most copper salts are insoluble
o copper hydroxide forms a precipitate (a solid)
o Balanced equation: CuCl2 (aq) + 2KOH (aq)  Cu(OH)2 (s) + 2KCl (aq)
• E.g2 what would happen if copper chloride and potassium hydroxide were mixed together?
o Sodium chloride + potassium carbonatesodium carbonate + potassium chloride
o All sodium and potassium salts are solubleno precipitate is formed (both products of the reaction are soluble)this is not a precipitation reaction
• Barium meals:
• In the diagnosis of intestinal problems, patients are made to swallow a drink - called a ‘barium meal’ - containing barium sulfate:
o As the barium sulfate passes through the patient’s digestive system, x-ray photos are taken
o Barium (like bone) absorbs x-rays shows up as white on the photos any problems with the patient’s digestive system can be seen
• Most barium salts are toxic, however…:
o barium sulphate is insoluble can’t enter the patient’s blood
o This makes it safe to swallow
ION TESTS
• Flame tests:
• Different metal ions produce different coloured flames when held over a Bunsen burner flame:
o Sodium (Na+) – yellow
o Potassium (K+) – lilac
o Calcium (Ca2+) – red
o Copper (II) (Cu2+) – green/blue
• Note: the most intense colours are obtained from solids, but flame tests also work when solids are dissolved in water (as aqueous solutions)
• Flame tests led to the discovery of new elements…:
o Chemists in the 1800s did flame tests of different samples of mineral water and then used a prism to separate the colours of light given off - method is called ‘spectroscopy’
o They saw a grey-blue colour that hadn’t been seen before and realised they had discovered a new element – later called it caesium
o A year later, using the same method, rubidium was discovered (gave off a dark red colour in a Bunsen flame)
• Precipitation tests:
• Some anions can be identified by precipitation tests…:
o Chloride ions (Cl-):
Add dilute nitric acid and silver nitrate to the solution
If the sample contains chloride ions, a white precipitate of silver chloride will form
o Sulfate ions (SO42-):
Add dilute hydrochloric acid and barium chloride to the solution
If solution contains sulfate ions, a white precipitate of barium sulfate will form
• Test for carbonate ions:
• Add a dilute acid to the solution
• If solution contains carbonate ions (CO3-), carbon dioxide gas will be given off, which when bubbled through limewater will turn the limewater milky

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Electron Configuration Diagrams | Properties of Matter | Chemistry | FuseSchool

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What are Isotopes? | Properties of Matter | Chemistry | FuseSchool

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The genius of Mendeleev's periodic table - Lou Serico

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Topic 1 – Atomic structure and the periodic table

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Topic 1 – Atomic structure and the periodic table

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#MENDELEEV
• In the 1800s, a chemist called Dmitri Mendeleev used the properties of elements known at the time to organise them into a table:
o Mendeleev arranged the elements in order of increasing atomic mass (i.e in order of increasing numbers of protons and neutrons – see below)
o Mendeleev organised elements with similar properties into vertical columns (called ‘groups’)
• Unlike other chemists before him, Mendeleev:
o sometimes broke the ‘increasing atomic mass rule’
e.g he switched tellurium and iodine around so that they would be in the same groups as elements with similar properties (i.e by switching them, iodine was next to bromine, chlorine, fluorine…)
o realised from the big jumps in atomic mass that there were still some elements to discover he left some gaps in his table
E.g he left two gaps between zinc and arsenic
Based on the known elements around them, Mendeleev predicted the properties of the elements which should go in the gaps
In the modern periodic table, these gaps have been filled by gallium and germanium – they have very similar properties to those predicted by Mendeleev
STRUCTURE OF THE ATOM
• Atoms are the smallest particles of an element that can take part in chemical reactions
• The 3 subatomic particles – protons, neutrons, electrons:
• At the centre of each atom is a nucleus containing protons and neutrons
• Note: the nucleus is about 20,000 times smaller than the overall size of the atom
• Electrons are arranged in shells (or ‘energy levels’) at different distances from the nucleus
• The masses and charges of subatomic particles are very small, they’re compared to those of a proton – these are called the ‘relative mass’ and the ‘relative charge’
Subatomic particle Relative mass Relative charge
Proton                     1                            +1
Neutron                     1                             0
Electron            Negligible (i.e 0)           - 1
• Atoms and elements:
• All atoms contain the same number of protons and electrons, atoms have no overall charge
• No two elements have the same number of protons in their atoms…e.g hydrogen has 1 proton, helium has 2 protons etc…
• Elements can have more than one atom…:
o Different atoms of the same element always have the same number of protons (the same number of electrons), but a different number of neutrons – these are called isotopes (see below)
• THE MODERN PERIODIC TABLE
• Atomic number – number of protons in the nucleus of an atom
• Mass number – total number of protons and neutrons in the nucleus of an atom
• In a chemical symbol, e.g  , the atomic number is the bottom number (the smaller one) and the mass number is the top number (the bigger one)
• From the symbol, we can calculate the number of protons, neutrons and electrons in an atom…e.g for  :
o Atomic number is 13 = 13 protons
o Mass number – atomic number = number of neutrons… 27 - 13 = 14 neutrons
o Number of protons = number of electrons = 13 electrons
• The elements in the modern periodic table are arranged in order of increasing atomic number rather than in order of increasing atomic mass as Mendeleev did
• The horizontal rows are called periods
• The vertical columns are called groups:
o Each group contains elements with similar properties
o The main groups are numbered 1-7 from left to right
o The group on the far right is called group 0
• Diagram of periodic table below:
o Elements to the left of the jagged line are metals
o Elements to the right of the line are non-metals
• Relative atomic mass (Ar):
• Atoms have very small masses = relative atomic mass is used instead of its actual mass in kilograms
• Relative atomic mass is the mass of an atom compared to that of carbon:
o Carbon has a relative atomic mass of 12
o Mass of a helium atom is one third that of carbon = its relative atomic mass is 4
• In the periodic table, each element has its atomic number (the bottom number) and its relative atomic mass (the top number) shown
Isotopes:
• Isotopes are different atoms of an element with the same number of protons and electrons, but different numbers of neutrons (i.e same atomic number, different mass number)
• E.g chlorine atoms always have 17 protons, but some can have 18 neutrons and others can have 20 neutrons – these are    and
• The presence of isotopes means that some relative atomic masses aren’t whole numbers…:
o Relative abundance (proportion) of Cl-35 is 75% (or 0.75)
o Relative abundance of Cl-37 is 25% (or 0.25)
o Relative atomic mass = (proportion in decimals x mass number) + (proportion in decimals x mass number)
o relative atomic mass of Cl = (0.75 x 35) + (0.25 x 37) = 35.5
• Relative atomic mass vs mass number:
• The mass number is the mass of an atom – i.e the number of protons and neutrons (because electrons have negligible relative mass)
• The relative atomic mass is the average mass of all the different atoms (isotopes) of an element (relative to carbon), taking their abundance into account
• In most cases, the relative atomic mass of an element is very similar to the mass number of the most common isotope (e.g in the above example, relative atomic mass of Cl is 35.5…mass number of most common isotope is 35)
ELECTRON SHELLS
• Electrons are arranged in shells around the nucleus of the atom
• Each shell is shown as a circle drawn around the chemical symbol for the atom
• The way in which electrons are arranged in an atom is called its electronic configuration
• Finding configurations:
• Different shells can contain different numbers of electrons…
• For the first 20 elements:
o The first shell can contain up to 2 electrons
o The second and third shells both hold up to 8 electrons
• The first shell fills up first, then the second shell, and so on…
• Elements which have the maximum number of electrons in their outer shells are said to have ‘full outer shells’
• Electronic configurations can be worked out using atomic numbers…e.g:
o Atomic number of sodium is 11= it has 11 protons = 11 electrons
o 2 electrons in first shell
o 8 electrons in second shell
o 1 electron in third (outer) shell
• This can be represented by a diagram:
• Its electronic configuration can also be written in the form ‘2.8.1’, whereby…:
o The numbers show how many electrons are in each shell
o The dots separate each shell
• Spotting connections:
• 1. The number of occupied shells is the same as the period number:
o E.g magnesium is in period 3 of the periodic table…its configuration is ‘2.8.2’…has 3 occupied shells
• 2. The number of outer electrons is the same as the group number (apart from elements in group 0 which all have full outer shells):
o E.g magnesium is in group 2 of the periodic table…its configuration is ‘2.8.2’…has 2 electrons in its outer shell