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  • Created on: 04-04-15 12:23

history of the periodic table

  • An English scientist called John Newlands put forward his Law of Octaves in 1864. He arranged all the elements known at the time into a table in order of relative atomic mass.
  • When he did this, he found a pattern among the early elements. The pattern showed that
    each element was similar to the element eight places ahead of it.
  • Newlands' table showed a repeating or periodic pattern of properties, but this pattern eventually broke down. By ordering strictly according to atomic mass, Newlands was forced to put some elements into groups which did not match their chemical properties.
  • For example, he put iron (Fe), which is a metal, in the same group as oxygen (0) and sulfur (S), which are two non-metals. As a result, his table was not accepted by other scientists.
  • In 1869, just five years after John Newlands put forward his Law of Octaves, a Russian chemist called Dmitri Mendeleev published a periodic table. Mendeleev also arranged the elements known at the time in order of relative atomic mass, but he did some other things that made his table much more successful.
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history of the periodic table

  • He realised that the physical and chemical properties of elements were related to their atomic mass in a 'periodic' way, and arranged them so that groups of elements with similar properties fell into vertical columns in his table.
  • Sometimes this method of arranging elements meant there were gaps in his horizontal rows or 'periods'. But instead of seeing this as a problem, Mendeleev thought it simply meant that the elements which belonged in the gaps had not yet been discovered.
  • He was also able to work out the atomic mass of the missing elements, and so predict their properties. And when they were discovered, Mendeleev turned out to be right.
  • For example, he predicted the properties of an undiscovered element that should fit below aluminium in his table. When this element, called gallium, was discovered in 1875, its properties were found to be close to Mendeleev's predictions. Two other predicted elements were later discovered, lending further credit to Mendeleev's table.
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the modern periodic table

  • However, the discovery of protons, neutrons and electrons in the early 20th century allowed Mendeleev’s table to be refined into the modern periodic table. It involved an important modification – the use of atomic number to order the elements.
  • An element’s atomic number (also called proton number) is the number of protons in its atoms.Using atomic number instead of atomic mass as the organising principle was first proposed by the British chemist Henry Moseley in 1913. It explained why Mendeleev needed to change the order of some of the elements in his table.
  • The elements in a vertical column are in the same group. The main groups are labelled groups 1-7, with the noble gases in group 0. All elements in a group have similar chemical properties.The elements in a group all have the same number of electrons in their highest occupied energy level (also referred to as the outer shell). This is why they have similar chemical properties.
  • An element’s group number is the same as the number of electrons in its highest occupied energy level (outer shell). For example, all the metals in Group 2 have 2 electrons in their highest occupied energy level (outer shell).Elements in a horizontal row are in the same period. The periods are numbered from top to bottom.
  • The period number is the same as the number of occupied energy levels (shells). For example, magnesium is in period 3 – its atoms have three occupied energy levels.
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alkali metals

  • All the Group 1 elements are very reactive. They must be stored under oil to keep air and water away from them. Group 1 elements form alkaline solutions when they react with water, which is why they are called alkali metals.
  • All the alkali metals react vigorously with cold water. In each reaction, hydrogen gas is given off and the metal hydroxide is produced. The speed and violence of the reaction increases as you go down the group. This shows that the reactivity of the alkali metals increases as you go down Group 1.
  • When lithium is added to water, lithium floats. It fizzes steadily and becomes smaller, until it eventually disappears.
  • When sodium is added to water, the sodium melts to form a ball that moves around on the surface. It fizzes rapidly, and the hydrogen produced may burn with an orange flame before the sodium disappears.
  • When potassium is added to water, the metal melts and floats. It moves around very quickly on the surface of the water. The hydrogen ignites instantly. The metal is also set on fire, with sparks and a lilac flame. There is sometimes a small explosion at the end of the reaction.
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The alkali metals

  • The hydroxides formed in all of these reactions dissolve in water to form alkaline solutions. These solutions turn universal indicator purple, showing they are strongly alkaline. Strong alkalis are corrosive.
  • The Group 1 elements have similar properties because of the electronic structure of their atoms - they all have one electron in their outer shell. in a reaction, an atom of a Group 1 element will form an ion with a single positive charge.

The reactivity of Group 1 elements increases as you go down the group because:

  • the atoms get larger as you go down the group
  • the outer electron gets further from the nucleus as you go down the group
  • the attraction between the nucleus and outer electron gets weaker as you go down the group so the electron is more easily lost
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halogens

properties and uses of halogens:

Element Properties Typical use

  • Chlorine Green gas sterilising water
  • Bromine Orange liquid Making pesticides and plastics
  • Iodine Grey solid Sterilising wounds
  • The halogens have low melting points and low boiling points. This is a typical property of non-metals. Fluorine has the lowest melting and boiling points. The melting and boiling points then increase as you go down the group.
  • Room temperature is usually taken as being 25°C. At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. There is therefore a trend in state from gas to liquid to solid as you go down the group.
  • The non-metal elements in Group 7 - known as the halogens - get less reactive as you go down the group. This is the opposite trend to that seen in the alkali metals in Group 1 of the periodic table.
  • Fluorine is the most reactive element of all in Group 7.
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Halogens

  • The reactivity of the halogens – the Group 7 elements - decreases as you move down the group. This can be shown by looking at displacement reactions.

-When chlorine (as a gas or dissolved in water) is added to sodium bromide solution, the chlorine takes the place of the bromine. Because chlorine is more reactive than bromine, it displaces bromine from sodium bromide.

  • The solution turns brown. This brown colour is the displaced bromine. The chlorine has gone to form sodium chloride.

This type of reaction happens with all the halogens. A more reactive halogen displaces a less reactive halogen from a solution of one of its salts.

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transition metals

Common properties:

  • The transition metals have these properties in common:
  • they are metals
  • they are good conductors of heat and electricity
  • they can be hammered or bent into shape easily
  • The transition metals are useful as construction materials. They are also useful for making objects that need to let electricity or heat travel through them easily.
  • Many transition metals act as catalysts in useful processes. For example, iron is the catalyst used catalyst in the Haber process when Making ammonia.
  • Many transition elements form ions with different charges. For example, iron forms iron(II) ions, Fe2+, and iron(III) ions, Fe3+. This means that iron oxide can exist in two forms, iron(II) oxide, FeO, and iron(III) oxide, Fe2O3.
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hardness of water

  • Rainwater is naturally weakly acidic because it contains carbonic acid, formed by the reaction between water and carbon dioxide in the atmosphere. When the rain falls, it flows over rocks or soaks into the ground and then passes through layers of rock. Compounds from the rocks dissolve into the water.
  • Hard water contains dissolved compounds, usually calcium or magnesium compounds. For example, limestone contains calcium carbonate, CaCO3. Carbonic acid in rainwater reacts with this to produce soluble calcium hydrogencarbonate:
  • carbonic acid + calcium carbonate → calcium hydrogencarbonate
  • H2CO3(aq) + CaCO3(s) → Ca(HCO3)2(aq)
  • The presence of calcium ions and magnesium ions in the water makes it hard. Soft water readily forms lather with soap, but it is more difficult to form lather with hard water. The dissolved calcium ions and magnesium ions in hard water react with the soap to form scum, so more soap is needed. Soapless detergents do not form scum with hard water.
  • The types of rocks found in different regions determines how hard or soft the water will be.
  • The water in some parts of the country is soft because it has low levels of dissolved calcium and magnesium compounds, while the water in other parts of the country is hard because it has higher levels of dissolved calcium and magnesium compounds.
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measuring hardness

Measuring hardness:

One way to measure the hardness in water is to perform a titration with soap solution:

  • A known volume of water is put into a conical flask. Soap solution is added to it from a burette or pipette. The mixture is swirled to help it form lather.
  • The volume of soap solution that needs to be added to form permanent lather is recorded. The harder the water, the greater the volume of soap solution needed.

Temporary hard water can be softened by boiling it. Permanent hard water stays hard, even when it is boiled.

  • Temporary hard water contains dissolved hydrogen carbonate ions, HCO3–. When heated, these ions decompose (break down) to form carbonate ions, CO32–. The carbonate ions in the boiled water react with dissolved calcium and magnesium ions to form insoluble precipitates (calcium carbonate and magnesium carbonate).
  • Permanent hard water contains dissolved sulfate ions, SO42–. These do not decompose when heated. They remain dissolved and do not react with calcium and magnesium ions - so the water stays hard even when boiled.
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benefits and drawbacks of hard water

Hard water has some benefits compared to soft water. For example, the dissolved calcium compounds in hard water:

  • can improve the taste of the water
  • are good for the development and maintenance of bones and teeth
  • can help to reduce heart disease
  • But hard water also has some drawbacks compared to soft water. For example:
  • More soap is needed to produce lather, which increases costs. This happens with temporary or permanent hardness.
  • The scum produced is unsightly - spoiling the appearance of baths and shower screens, for example.
  • Temporary hardness can reduce the efficiency of kettles and heating systems. This is because limescale (a solid containing calcium carbonate) is produced when the water is heated. It coats the heating element in kettles, and the inside of boilers and hot water pipes.
  • This means more energy is needed to heat the water, again increasing costs. Pipes may become blocked by limescale - causing the heating system to break down.
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removing water hardness

  • Sodium carbonate, Na2CO3, is also known as washing soda. It can remove temporary and permanent hardness from water. Sodium carbonate is soluble but calcium carbonate and magnesium carbonate are insoluble. The carbonate ions from sodium carbonate react with the calcium and magnesium ions in the water to produce insoluble precipitates. For example:
  • calcium ions + sodium carbonate → calcium carbonate + sodium ions
  • Ca2+(aq) + Na2CO3(aq) → CaCO3(s) + 2Na+(aq)
  • The water is softened because it no longer contains dissolved calcium ions and magnesium ions. It will form lather more easily with soap. However, the calcium carbonate and magnesium carbonate precipitates to form limescale. As well as being unsightly on your taps, it can also clog up pipes in heating systems (causing them to break down). This makes treatment with sodium carbonate suitable for softening water only in certain circumstances - such as softening water for hand washing clothes.
  • Commercial water softeners often use ion exchange resins. These substances are usually made into beads, which are packed into cylinders called ion exchange columns. These can be built into machines, such as dishwashers, or plumbed into water systems to continuously soften the water.The resin beads have sodium ions attached to them. As the hard water passes through the column, the calcium and magnesium ions swap places with the sodium ions.
  • The calcium and magnesium ions are left attached to the beads, while the water leaving the column contains more sodium ions. The hard water is softened because it no longer contains calcium or magnesium ions.
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water quality

  • Solids in the water, such as leaves and soil, must be removed. The water is sprayed onto specially-prepared layers of sand and gravel called filter beds.
  • Different-sized insoluble solids are removed as the water trickles through the filter beds. These are cleaned every so often by pumping clean water backwards through the filter.
  • The water is then passed into a sedimentation tank. Aluminium sulfate is added to clump tiny particles together to make larger particles, which settle out more easily. The water is then passed through a fine filter, such as carbon granules, to remove very small particles.
  • Chlorine is added to drinking water to sterilise it. The chlorine kills microbes - including microbes that cause potentially-fatal diseases such as typhoid, cholera and dysentery.

Adding fluoride to the water supply:

  • Results from scientific research indicate that fluoridated water can improve dental health by reducing tooth decay. Many areas of the country naturally have low fluoride levels present in the water supply.
  • Some people argue that extra fluoride should not be added to water, even if it does improve dental health. They claim that fluoridation:
  • has been linked to tooth mottling (staining), bone disease and pain
  • forces people to consume fluoride when they drink tap water - taking away their choice.
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water quality

  • Solids in the water, such as leaves and soil, must be removed. The water is sprayed onto specially-prepared layers of sand and gravel called filter beds.
  • Different-sized insoluble solids are removed as the water trickles through the filter beds. These are cleaned every so often by pumping clean water backwards through the filter.
  • The water is then passed into a sedimentation tank. Aluminium sulfate is added to clump tiny particles together to make larger particles, which settle out more easily. The water is then passed through a fine filter, such as carbon granules, to remove very small particles.
  • Chlorine is added to drinking water to sterilise it. The chlorine kills microbes - including microbes that cause potentially-fatal diseases such as typhoid, cholera and dysentery.

Adding fluoride to the water supply:

  • Results from scientific research indicate that fluoridated water can improve dental health by reducing tooth decay. Many areas of the country naturally have low fluoride levels present in the water supply.
  • Some people argue that extra fluoride should not be added to water, even if it does improve dental health. They claim that fluoridation:
  • has been linked to tooth mottling (staining), bone disease and pain
  • forces people to consume fluoride when they drink tap water - taking away their personal choice (making it unethical).
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water filtering

  • Some people prefer to filter their water rather than use it straight from the tap. Filtering removes impurities and this can improve the taste and quality of the water. Filtering also helps to soften the water.
  • Commercially-available systems use cartridges containing water filters. These may be fitted in jugs or kettles, or plumbed in to the home water supply pipework.

The filter cartridges can contain:

  • silver to kill bacteria
  • carbon (‘activated charcoal’) to absorb impurities, eg chlorine
  • ion exchange resins to soften the water, and remove heavy metal ions (such as lead ions)
  • Silver nanoparticles have an antibacterial effect. Their presence in the filter prevents the growth of bacteria within the filter if water is left inside it for long periods. Silver nanoparticles also help break down harmful pesticides which might be in the water.
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reversible reactions

  • When reversible reactions reach equilibrium the forward and reverse reactions are still happening but at the same rate, so the concentrations of reactants and products do not change. The balance point can be affected by temperature, and also by pressure for gasses in equilibrium.

Chemical equilibrium

  • In a chemical equilibrium, the concentrations of reactants and products do not change. But the forward and reverse reactions have not stopped - they are still going on at the same rate as each other.
  • Imagine walking the wrong way on an escalator - at the same speed as the escalator, but in the opposite direction. Your legs would still be walking forwards, and the escalator would continue to move backwards. However, the net result would be that you stay in exactly the same place. This is what happens in an equilibrium.

Other factors

  • If we remove the products from an equilibrium mixture, more reactants are converted into products. If a catalyst is used, the reaction reaches equilibrium much sooner, because the catalyst speeds up the forward and reverse reactions by the same amount. The concentration of reactants and products is nevertheless the same at equilibrium as it would be without the catalyst.
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reversible reactions

  • If the forward reaction is exothermic and the temperature is increased, the yield of products is decreased. If the temperature is decreased, the yield of products is increased.
  • If the forward reaction is endothermic and the temperature is increased, the yield of products is increased. If the temperature is decreased, the yield of products is decreased.
  • Changing the pressure has little effect on an equilibrium mixture without gases - but can have a big effect on an equilibrium mixture containing gases. If the pressure is increased, the position of equilibrium moves in the direction of the fewest molecules.
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the harbour process

The raw materials for this process are hydrogen and nitrogen:

hydrogen is obtained by reacting natural gas with steam, or from cracking oil fractions

nitrogen is obtained from the air. Air is 78 per cent nitrogen and nearly all the rest is oxygen. When hydrogen is burned in air, the oxygen combines with the hydrogen, leaving nitrogen behind.

In the Haber process, nitrogen and hydrogen react together under these conditions:

a high temperature - about 450ºC, a high pressure, an iron catalyst.

  • In addition, any unreacted nitrogen and hydrogen are recycled.The reaction is reversible

Different factors affect the cost of making a new substance. Factors that increase cost include:

  • high pressures (they increase the cost of the equipment)
  • high temperatures (they increase the energy costs).

Factors that decrease:

  • catalysts (they increase the rate of reaction)
  • recycling unreacted starting materials
  • automating equipment (because fewer people need to be employed, cutting the wage bill).
  • The pressure chosen for the Haber process is a compromise. A high pressure increases the percentage yield of ammonia but very high pressures are expensive.
  • The temperature chosen is also a compromise. A high temperature gives a fast reaction but decreases the percentage yield of ammonia. 450°C gives a reasonably fast reaction with a sufficiently high percentage yield of ammonia.
  • The use of an iron catalyst increases the rate of the reaction, but it does not alter the percentage yield of ammonia.
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alcohols

  • alcohols all contain the functional group –OH. This group is responsible for the properties of alcohols.
  • The names of alcohols end with ‘-ol’ – eg ethanol.
  • The first three alcohols in the homologous series are methanol, ethanol and propanol. These alcohols are highly flammable, making them useful as fuels. They are also used as solvents in marker pens, medicines, and cosmetics (such as deodorants and perfumes).
  • Ethanol is the alcohol found in alcoholic drinks such as wine and beer. Ethanol is usually mixed with petrol for use as a fuel (see the Biology revision bite on Biofuels).

Alcohol Number of carbon atoms Structural formula Displayed formula

Methanol 1 CH3OH

Ethanol 2 CH3CH2OH

Propanol 3 CH3CH2CH2OH

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alcohols

The alcohols methanol, ethanol and propanol all have the following properties:

-They are colourless liquids that dissolve in water to form a neutral solution (pH7).

  • They react with sodium to produce hydrogen and a salt. For example:
  • ethanol + sodium = hydrogen + sodium ethoxide

This reaction is similar but less vigorous to the reaction of water with sodium. This is due to the similarity in structure between water and the –OH group in alcohols. They burn in the air, releasing energy and producing carbon dioxide and water.

Balanced equations for alcohols
methanol + oxygen = carbon dioxide + water

2CH3OH(l) + 3O2(g) = 2CO2(g) + 4H2O(l)

ethanol + oxygen = carbon dioxide + water

2C2H5OH(l) + 6O2(g) = 4CO2(g) + 6H2O(l)

propanol + oxygen = carbon dioxide + water

2C3H7OH(l) + 9O2(g) = 6CO2(g) + 8H2O(l)

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carboxylic acids

carboxylic acids all contain the same functional group –COOH.

The names of carboxylic acids end in ‘-oic acid’ – eg ethanoic acid.

Carboxylic acid Number of C atoms Structural formula Displayed formula
Methanoic acid 1 HCOOH
Ethanoic acid 2 CH3COOH
Propanoic acid 3 CH3CH2COOH

Ethanoic acid from ethanol
Vinegar is an aqueous solution containing ethanoic acid. Ethanoic acid is formed from the mild oxidation of the ethanol (which is an alcohol). This can be achieved through:

The addition of chemical oxidising agents - such as acidified potassium dichromate.

The action of microbes in aerobic conditions (in the presence of oxygen). This happens on a small scale when a bottle of wine is left open and exposed to air. On a commercial scale, it is achieved in a fermenter using acetic acid bacteria.

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carboxylic acids

Carboxylic acids have the following properties:

They dissolve in water to produce acidic solutions (pH less than 7). They react with carbonates to produce carbon dioxide and a salt and water. For example:

calcium carbonate + ethanoic acid → calcium ethanoate + water + carbon dioxide

They all react with alcohols, in the presence of an acid catalyst, to form esters. For example:

ethanol + ethanoic acid → ethyl ethanoate + water

Strong acid, such as hydrochloric acid, ionise fully in water:

HCl(aq) → H+(aq) + Cl–(aq)

Their aqueous solutions have a high concentration of hydrogen ions, H+. This gives them a low pH. Carboxylic acids are weak acids. They do not completely ionise when they are dissolved in water. Instead only some of their molecules ionise to form H+ ions:

CH3COOH(aq) CH3COO–(aq) + H+(aq)

This means that an aqueous solution of a weak acid will have a higher pH compared to the same concentration of an aqueous solution of a strong acid.

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esters

  • Esters are a group of organic compounds which all contain the functional group –COO–. They have these properties in common:
  • they are volatile - they are liquids that become vapours easily
  • they have distinctive ‘fruity’ smells
  • These properties make esters very useful as food flavourings, and as perfumes in cosmetics. Some esters are obtained from natural sources, such as fruits. Others are manufactured.

Making ethyl ethanoate:

  • Ethyl ethanoate is the ester made from ethanol and ethanoic acid. Sulfuric acid is added to act as a catalyst in the reaction.

ethanol + ethanoic acid ethyl ethanoate + water
CH3CH2OH(aq) + CH3COOH(aq) CH3CH2OOCCH3(aq) + H2O(l)

  • The distinctive smell of ethyl ethanoate (which is like modelling glue) can be detected as the reaction proceeds. Excess ethanoic acid in the reaction mixture is neutralised with sodium hydrogencarbonate, then a few drops of the mixture added to water so that the smell can be detected more effectively.
  • The first part of an ester’s name comes from the alcohol - it ends with the letters 'yl'. The second part of its name comes from the carboxylic acid - it ends with the letters 'oate'.
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moles and titration

  • The concentration of an acid or alkali can be calculated by carrying out an experiment called a titration. You should be able to identify the apparatus needed to carry out a simple acid-alkali titration, and to describe how it is done.

Materials

The apparatus needed includes a:

  • pipette to accurately measure a certain volume of acid or alkali
  • pipette filler to use the pipette safely
  • conical flask to contain the liquid from the pipette
  • burette to add small, measured volumes of one reactant to the other reactant in the conical flask
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titration

  • Use the pipette and pipette filler to add 25 cm3 of alkali to a clean conical flask.
  • Add a few drops of indicator and put the conical flask on a white tile (so you can see the colour of the indicator more easily).
  • Fill the burette with acid and note the starting volume.
  • Slowly add the acid from the burette to the alkali in the conical flask, swirling to mix.
  • Stop adding the acid when the end-point is reached (the appropriate colour change in the indicator happens). Note the final volume reading.

The same method works for adding an alkali to an acid - just swap around the liquids that go into the conical flask and burette.

The titre:
The difference between the reading at the start and the final reading gives the volume of acid (or alkali) added. This volume is called the titre.

For example, if the reading at the start is 1.0 cm3 and the final reading is 26.5 cm3, then the titre is 25.5 cm3 (26.5 – 1.0). Note that the titre will depend upon the volume of liquid in the conical flask, and the concentrations of the acid and alkali used.

It is important to repeat the titration several times to check that your titre value is consistent so that your calculations are reliable.

If universal indicator is used, the colour changes gradually through a range of colours. On the other hand, a single indicator like litmus or phenolphthalein gives a sharp end-point where the colour changes suddenly.

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energy

  • Heat energy can be given out or taken in from the surroundings during chemical reactions. The amount of energy transferred can be measured. This is called calorimetry.

Energy changes from combustion:

  • The diagram shows a simple calorimetry experiment to measure the heat energy released from burning a fuel. You should be able to recognise and label apparatus like this.
  • The spirit burner containing the fuel is usually weighed before and after the experiment - in this way, the mass of the fuel burned can be found. Knowing the mass of fuel burnt and the temperature change in the water, it is then possible to calculate the energy released by the fuel. This method also works for finding the amount of energy released by foods.
  • The biggest source of error is usually heat loss to the surroundings. This can be reduced by insulating the sides of the calorimeter and adding a lid.

Energy changes from reactions in solution:

Energy changes also happen when chemicals in solution react. For example, heat energy is given out to the surroundings when acids and alkalis react together.

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energy

During a chemical reaction:

  • bonds in the reactants are broken
  • new bonds are made in the products
  • Energy is needed to break bonds, and energy is released when bonds are made.

Exothermic reactions:

  • Exothermic reactions give out heat energy to the surroundings. Exothermic reactions have a negative energy change. Energy released in an exothermic reaction

Some examples of exothermic reactions are:

  • combustion
  • neutralisation reactions between acids and alkalis
  • the reaction between water and calcium oxide

Endothermic reactions:

  • Endothermic reactions absorb heat energy from the surroundings, making the temperature of the surroundings cooler. Endothermic reactions have a positive energy change. This is shown in the energy level diagram below. Energy is absorbed in an endothermic reaction

Some examples of endothermic reactions are:

  • electrolysis
  • the reaction between ethanoic acid and sodium carbonate
  • the thermal decomposition of calcium carbonate in a blast furnace

Bond breaking and making:
-In an exothermic reaction, more energy is released when new bonds are made than is needed to break existing bonds.

  • In an endothermic reaction, more energy is needed to break existing bonds than is released when new bonds are made.
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energy

  • in an exothermic reaction the energy level of the reactants is higher than the energy level of the products.
  • However, the energy curve goes up from the reactants’ energy level to begin with, then drops to the products’ energy level. This is because many reactions need an input of energy to start the reaction off. This is energy is called the activation energy. It is represented on an energy level diagram as the difference between the reactants’ energy level and the top of the curve.

For example, burning methane in a Bunsen burner:

  • methane + oxygen → carbon dioxide + water
    CH4 + 2O2 → CO2 + 2H22O
  • A catalyst is a substance that speeds up the rate of a chemical reaction without being used up in the reaction.
  • Catalysts can do this because they provide a different pathway for the reaction to follow. This pathway has lower activation energy than the one followed by the uncatalysed reaction. As a result, a greater proportion of reacting particles have enough energy to react.

Energy changes in chemical reactions

  • Lowering the activation energy has many advantages. It means that reactions happen more quickly and are more economical in terms of the energy required for industrial-scale reactions.
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getting energy from hydrogen

Hydrogen is often seen as an environmentally-friendly alternative to fossil fuels. Some car manufacturers have developed cars than run on hydrogen rather than petrol or diesel.

There are two ways in which hydrogen is used to power cars:
Burning hydrogen directly in the engine, Water is the only product formed when hydrogen burns:

  • hydrogen + oxygen → water
  • 2H2 + O2 → 2H2O
  • There are no carbon dioxide emissions that could contribute to global warming.

Hydrogen fuel cells

  • In a hydrogen fuel cell, hydrogen reacts with oxygen without burning. The energy released is used to generate electricity, which is used to drive an electric motor.
  • Problems with hydrogen
    At the moment, most hydrogen is made by reacting steam with coal or natural gas - both non-renewable resources.
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advantages and disadvantages of hydrogen

  • Hydrogen can also be made by passing electricity through water. Unfortunately, most electricity is generated using coal and other fossil fuels, so pollution from burning these fuels happens at the power station. Pollution therefore still occurs.
  • However, some countries are producing hydrogen using electricity from renewable sources, such as geothermal energy in Iceland.

Advantages and disadvantages of hydrogen

There are some benefits to using hydrogen as a fuel:

  • unlike petrol and diesel, hydrogen does not generate carbon dioxide when burnt
  • hydrogen fuel cells are very efficient
  • However, there are also some downsides too:
  • few filling stations sell hydrogen
  • hydrogen must be compressed and liquefied, and then stored in tough, insulated fuel tanks
  • atmospheric pollution may be generated during the production of hydrogen
  • hydrogen fuel cells do not work at very low temperatures, and they may also require a platinum catalyst (platinum is expensive and prone to contamination by impurities)
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test for positive ions

  • Metals change the colour of a flame when they are heated in it. Different metals give different colours to the flame, so flame tests can be used to identify the presence of a particular metal in a sample.
    This is how you would carry out a typical flame test:
  • Dip a clean flame test loop in the sample solution
  • Hold the flame test loop at the edge of a Bunsen burner flame
  • Observe the changed colour of the flame, and decide which metal it indicates
  • Clean the loop in acid and rinse with water, then repeat steps 1 to 3 with a new sample

Some common metals and their flame test colours
Metal Flame test colour
Barium Pale green
Calcium Yellow-red
Copper Green-blue
Lithium Red
Sodium Orange
Potassium Lilac

for example: Flame tests are useful for confirming the results of a precipitate test. For example, an unknown solution that produced a pale blue precipitate with sodium hydroxide solution, and a green-blue flame test, must contain a copper compound.
To identify an alkali metal, a flame test must be used instead of a sodium hydroxide precipitate test. This is because the alkali metals do not form precipitates with sodium hydroxide.

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precipitate tests

Transition metals form coloured compounds when they react with other elements. Many of these are soluble in water and so form coloured solutions. Others are not soluble and so form precipitates (the insoluble products of these reactions).

  • To test for these metal ions, sodium hydroxide solution is added to them.
    Here is the equation for copper sulfate solution reacting with sodium hydroxide solution:
  • copper sulphate + sodium hydroxide → copper hydroxide + sodium sulphate
  • CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4
  • (blue solution + colourless solution → blue precipitate + colourless solution)

Different transition metals form different coloured precipitates.

metal Colour of precipitate

Al3+ aluminium: white

Ca2+ calcium: white

Cu2+ copper: blue

Fe2+ iron(II): green

Fe3+ iron(III): brown

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test for negative ions

Testing for carbonate ions:

  • Carbonate ions, CO32- can be detected whether in a solid compound or in solution. An acid, such as dilute hydrochloric acid, is added to the test compound.
  • Carbon dioxide gas bubbles if carbonate ions are present. Limewater is used to confirm that the gas is carbon dioxide. It turns from clear to milky when carbon dioxide is bubbled through.

Testing for sulfate ions

  • Sulfate ions in solution, SO42-, are detected using barium chloride solution. The test solution is acidified using a few drops of dilute hydrochloric acid, and then a few drops of barium chloride solution are added. A white precipitate of barium sulfate forms if sulfate ions are present.

For example:

barium chloride + sodium sulfate = sodium chloride + barium sulfate

BaCl2(aq) + Na2SO4(aq) = 2NaCl(aq) + BaSO4(s)

  • The hydrochloric acid is added first to remove any carbonate ions that might be present - they would also produce a white precipitate, giving a false positive result.
  • Barium nitrate solution can be used instead of barium chloride solution. However, nitric acid is added first to acidify the test solution. Sulfuric acid cannot be used because it contains sulfate ions - these would interfere with the second part of the test.
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tests for negative ions

  • The halogens are the elements in Group 7 of the periodic table. Chlorine, bromine and iodine are halogens. Their ions are called halide ions, eg chloride, Cl-. Halide ions in solutions are detected using silver nitrate solutions. The test solution is acidified using a few drops of dilute nitric acid, and then a few drops of silver nitrate solution are added.

Different coloured silver halide precipitates form, depending on the halide ions present:

  • chloride ions give a white precipitate of silver chloride
  • bromide ions give a cream precipitate of silver bromide
  • iodide ions give a yellow precipitate of silver iodide

silver nitrate + sodium bromide = sodium nitrate + silver bromide
AgNO3(aq) + NaBr(aq) = NaNO3(aq) + AgBr(s)

One way to remember the colours is to think of ‘milk, cream, butter’ (white, cream, yellow).

  • The nitric acid is added first to remove any carbonate ions that might be present - they would produce a white precipitate of silver carbonate, giving a false positive result for chloride ions.

Testing for nitrate ions
Nitrate ions (NO3-) can be detected by reducing them to ammonia. This is done by:

  • adding sodium hydroxide solution, then aluminium powder or foil
  • heating strongly,
  • If nitrate ions are present, ammonia gas is given off. This has a characteristic choking smell. It also turns damp red litmus paper or damp universal indicator paper blue.
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