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    Reaction types Reaction types Presentation Transcript

    • Types of Reactions Yr 10 Chemistry
      • There are many different types of reaction in chemistry
      • We are going to look at a selection
      Can you name any?
    • Reactions
      • Dissociation
      • Ionisation
      • Combustion
      • Precipitation
      • Decomposition
      • Displacement
      • Neutralisation
      • Redox
    • Reaction Rates
      • Every different chemical reaction occurs at a different speed, or a different rate.
      • The study of the rates of reactions is often referred to as reaction kinetics.
    • Rates of Reaction Chemical reactions occur when different atoms or molecules collide:
      • For the reaction to happen the particles must have a certain amount of energy – this is called the ACTIVATION ENERGY.
      • The rate at which the reaction happens depends on many things:
      • The temperature of the reactants,
      • Their concentration
      • Their surface area
      • The reactivity of the reactants
      • Whether or not a catalyst is used
    • Rate of reaction graph Amount of product formed Time Slower reaction Fast rate of reaction here Slower rate of reaction here due to reactants being used up
    • Reaction Rates
      • Reactions can occur
      • VERY fast eg explosions
      • Very slow eg. Coal converting to diamonds.
      • Or anywhere in between.
      • The rate is the speed of conversion from the reactants (the things that are combining) to the products (the things produced).
      • We will look at the factors influencing reaction rates in isolation, not as a entire system.
    • Reaction Rates Collision Theory
      • Reactions occur when 2 reactive molecules collide with each other.
      • The rate at which these collisions occur governs the reaction rate.
      • These collisions must occur with sufficient energy and at the appropriate orientation.
    • Reaction Rate Effects
      • Temperature
        • As the temperature increases, the rate of reaction increases.
        • This is because the particles are moving more rapidly when they are heated, therefore they are much more likely to collide and thus react with each other.
        • The collisions are much more likely to have sufficient energy to allow a reaction to occur.
        • As a rule of thumb a reaction rate increases about 10 fold for each 10 o C rise in temperature.
        • Many reactions give off heat (exothermic), therefore they will cause an acceleration of the reaction as they proceed.
    • Reaction Rate Effects
      • CONCENTRATION:
        • An increase in concentration will bring about an increase in the reaction rate as if there are more particles then they are more likely to collide.
    • Concentration
      • If we make one reactant more concentrated (like making a drink of orange squash more concentrated )
      • There are more particles in the same volume to react
      • So the reaction goes faster.
      There are less red particles in the same volume so there is less chance of a collision There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster
    • Reaction Rate Effects
      • SURFACE AREA:
        • An increase in the surface area increases the likelihood of a collision as there is a greater area on which those collisions can occur.
        • A reduction in particle size will bring about an increase in surface area. Therefore there is a greater area on which the reaction can occur.
        • Particles that have been ground up into a powder are much more reactive than when in a solid form. This is why many dusts such as wheat dust or coal dust are highly explosive.
    • Surface area
      • If we make the pieces of the reactants smaller we increase the number of particles on the surface which can react.
      • This makes the reaction faster.
      The particles on the surface can react When cut into smaller pieces there are more particles on the outside that can react
    • Reaction Rate Effects
      • REACTIVITY:
        • In general terms, excluding the noble gases, the most reactive elements are those on the sides of the periodic table. (Gp I and Gp VII)
    • Measuring rate of reaction Two common ways: 1) Measure how fast the products are formed 2) Measure how fast the reactants are used up
    • Activation Energy
      • There is a minimum amount of energy required for reaction: the activation energy .
      • Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier.
    • Exothermic v Endothermic
      • Energy is put into the system to get it “over the activation energy hill”
      • If there is a lot of energy out, the reaction is exothermic and the system rolls down the energy hill to a lower point.
      • If the energy out is less than the activation energy, the reaction is endothermic
    • Activation Energy
      • This can be easily understood when the ignition of flammable liquids is considered.
      • Compare how easy it is to light a series of flammable liquids such as methanol, ethanol, hexane and ether.
      • Also compare this to your knowledge of igniting fuels such as petrol and diesel.
    • Catalysts
      • A catalyst is something that causes a process to proceed at a faster rate without being itself consumed by the process.
      • An example of this is the platinum inside a car’s exhaust system which allows the conversion of exhaust gases to a more environmentally friendly form. The platinum facilitates the reaction, but does not get used up by the process.
    • Catalytic Converter
    • Catalysts
      • Catalysts can be adversely affected by some substances which can inhibit their performance.
      • Such materials are referred to as poisons. An example of this is the presence of lead in petrol. A contamination of lead in fuel poisons the platinum catalytic converter and permanently prevents it from operating.
    • Catalysts
      • Many life processes are facilitated by natural catalysts called enzymes. Without these catalysts, the activation energy is too high to allow rapid reactions, and many bodily functions would cease.
      • Without enzymes in the stomach, the process of digestion would be extremely slow, and most of the nutritional value of the foods eaten would be excreted.
    • Dissociation Yr 10 Chemistry
    • Dissociation reactions
      • Are reactions that involve setting free ions
      • They occur when an ionic solid dissolves in water.
      • The positive and negative ions in the lattice separate in solution and are now free to move around
    • Example
      • Sodium chloride dissociates when added to water
      • NaCl (s) + aq
      • Na + (aq) + Cl - (aq)
    • Solubility
      • In order for dissociation to happen the ionic compound must be SOLUBLE in water.
      • The table shows soluble anions, in addition all compounds of potassium, sodium and ammonium are soluble
      Anion Formula exceptions Acetate CH 3 COO - Nitrate NO 3 - Chloride Cl - Silver, mercury, lead Sulfate SO 4 2- Barium, strontium, lead
    • Questions
      • For each compound record either a null result or write the dissociation equation
      • Sodium sulfate calcium hydroxide
      • Lithium nitrate aluminium chloride
      • Potassium hydroxide barium sulfate
      • Copper sulfide ammonium carbonate
    • Solutions
      • Sodium sulfate Na + + SO 4 2-
      • calcium hydroxide null
      • Lithium nitrate Li + + NO 3 -
      • aluminium chloride Al 3+ + Cl -
      • Potassium hydroxide K + + OH -
      • barium sulfate null
      • Copper sulfide null
      • ammonium carbonate NH 4 + + CO 3 2-
    • Ionisation Yr 10 Chemistry
    • Ionisation reactions
      • Are reactions where ions are formed
      • They occur when a molecule dissolves in water and a covalent bond is broken creating ions.
    • Example
      • Hydrogen chloride is a covalent compound which separates into H + and Cl - dissolved in water
      • HCl (g) + H 2 O (l) H + (aq) + Cl - (aq)
      • Ammonia is another covalent compound that separates when dissolved in water
      • NH 3(g) + H 2 O (l) NH 4 + (aq) + OH - (aq)
    • Are these ionisation reactions?
      • K 2 SO 4 K + + SO 4 2-
      • SO 2 + H 2 O H 2 SO 3
      • PbCl 2 Pb 2+ + Cl -
      • HF+H 2 O H 3 O + + F -
      • NaOH Na + + OH -
    • Are these ionisation reactions?
      • Nitrous oxide and water
      • Calcium carbonate and water
      • Lead oxide and water
      • Ammonium nitrate and water
      • Carbon dioxide and water
    • Combustion Yr 10 Chemistry
    • Combustion
      • Substances can burn in oxygen to form an oxide
      • We will consider metals and hydrocarbons
    • Metal + Oxygen
      • Many metals react with oxygen to form oxides of the metal.
      • The standard reaction is:
        • metal + oxygen  metal oxide
      • eg:
        • Iron + oxygen  iron oxide
        • Magnesium + oxygen  magnesium oxide
    • Metal + Oxygen Magnesium in Oxygen
    • Hydrocarbon + Oxygen
      • Hydrocarbons are compounds made of only hydrogen and carbon
      • They burn in oxide to form the oxides of their components
        • Carbon dioxide
        • Dihydrogen monoxide (water)
    • Hydrocarbon + Oxygen
      • Eg
      • Methane + oxygen carbon dioxide + water
      • Propane
      • Butane
      • Octane
      • Now can you write them in symbol form?
      Can you write equations for the combustion of
    • Hydrocarbon + Oxygen
      • Propane + oxygen carbon dioxide + water
      • Butane + oxygen carbon dioxide + water
      • Octane + oxygen carbon dioxide + water
    • Precipitation Yr 10 Chemistry
    • Precipitation Reactions
      • Are any reaction in aqueous form that results in an insoluble product
      • You need to remember solubilities to predict these reactions
      • Precipitation is the reverse of dissolution
    • Solubility
      • Salts are substances composed of ions; many salts are soluble in water
      • Solubility is the amount of a substance needed to make a saturated solution at a specific temperature.
      • Solubility of solids and liquids varies widely, for example:
      • 35.7g of NaCl dissolves in 100mL of H 2 O, but only 0.044g of PbI 2 dissolves in 100mL of H 2 O
    • Solubility anion formula Solubility Acetate C2H3O2 - All soluble Nitrate NO3 - All soluble Chloride Cl - Not Ag, Hg, Pb Sulfate SO4 2- Not Ba, Sr, Pb Carbonate CO3 2- Insoluble* Hydroxide OH - Insoluble* Sulfide S 2- Insoluble Not group I or II Na/K/NH4 salts All soluble
    • Solubility
      • Solubility is temperature dependent.
      • For gases, an increase in temperature will decrease the amount of gas that can dissolve
      • For salts, the amount and rate of dissolving will increase as temperature increases :
      Temperature Amount & Rate of Dissolving of Salt
    • Precipitate Formation: Formation of Lead (II) Iodide When mixed, two clear solutions produce a dense yellow precipitate To observe the formation of a precipitate from ionic compounds in an aqueous state.
    • Demonstration
      • The formation of PbI 2(s) from Pb(NO 3 ) 2 and 2KI solutions.
      Pb 2- NO 3 - K + I - + PbI 2(s) K + NO 3 -
    • Double Replacement Reactions
      • Occurs when the elements in a solution of reacting compounds exchange places, or replace one another.
      • AB + CD  AD + CB
      • For this demonstration:
      • Pb(NO 3 ) 2(aq) + 2KI (aq)  PbI 2(s) + 2KNO 3(aq)
    • Precipitates
      • A precipitate is an insoluble solid formed when two ionic solutions react.
      • Pb(NO 3 ) 2 and KI are ionic compounds that dissociate into ions as they dissolve in water, but when mixed together, undergo a double replacement reaction to form an insoluble yellow precipitate, PbI 2 .
    • Write equations for …
      • Sodium chloride + silver nitrate
      • Lead sulfide + copper chloride
      • Barium chloride + potassium sulfate
      • Lead nitrate + sodium carbonate
      • Magnesium chloride + potassium carbonate
      • Barium nitrate + potassium sulfate
    • Practical
      • Carry out precipitation practical
    • Decomposition Yr 10 Chemistry
    • Decomposition reactions
      • Are those where a compound breaks down when heated
      • Calcium carbonate decomposes with heat into calcium oxide and carbon dioxide
      • CaCO 3 CaO + CO 2
      • Remember carbon dioxide can be tested for using limewater
    • Practical
      • You will decompose calcium carbonate (white solid) to form calcium oxide (white solid) and carbon dioxide (colourless gas)
      • How will you know the reaction has occurred?
      • What test could you do?
    • Which decomposes quickest?
      • The following carbonate solutions are provided
      • Sodium, calcium, magnesium, copper, zinc
      • Your challenge is to design a practical that will allow you to find out which decomposes quickest on heating
    • Displacement Reactions Year 10 Chemistry
    • The Activity Series of Metals
      • Remember that metals vary in their activity.
      • Chemists have made a list and placed the metals in order of their activity.
    • The Activity Series of Metals most reactive least reactive K Ca Na Mg Al Zn Fe Cu Ag Au
    • Displacement Reactions
      • Because some metals are more active than others, they want to react and form compounds more than others.
      • If a more reactive metal is combined with a less reactive metal salt, the more reactive metal will displace the less reactive metal from its salt.
      • The less reactive metal will precipitate out of the solution.
    • Displacement Reactions
      • For example consider what happens if Mg is mixed with a solution of CuSO 4 .
      • Mg + CuSO 4 
      • Magnesium is more active than copper.
      • The magnesium pushes the copper out of solution.
      • The magnesium reacts with the SO 4 ion forming magnesium sulfate.
      • Copper metal forms.
      Mg SO 4 + Cu (s)
    • Displacement Reactions
      • Consider :
      • Fe + Al 2 (SO 4 ) 3 
      • In this case, according to the activity series, aluminium is the more active metal.
      • It has already reacted with the sulfate group.
      • The iron is not active enough to push the aluminium out of solution.
      • No visible reaction is observed.
      NVR
    • Is it displacement?
      • If it is displacement then write out the equation
      • Na + CuS
      • Fe + Ag 2 O
      • CuSO 4 + Mg
      • MgO + K
      • Na 2 O + Al
      • Al 2 O 3 + Ca
    • Neutralisation – Reactions with Acids Year 10 Chemistry
    • Reactions of Acids
      • There are some standard reactions of acids:
        • acid + alkali
        • acid + metal
        • acid + metal oxide
        • acid + carbonate
      • These reactions are called Neutralisation Reactions because they neutralise an acid.
    • PROPERTIES OF ACIDS & BASES
      • Acids
        • Produce hydrogen ions (H + ) in H 2 O
        • Taste sour
        • Turn blue litmus (lichen)  red
        • Act as electrolytes in solution
        • Neutralise solutions containing hydroxide ions (OH - )
        • React with several metals releasing H 2(g)  corrosion
        • React with carbonates releasing CO 2(g)
        • Destroy body tissue
      • Bases
        • Produce or cause an increase in hydroxide ions (OH - ) in H 2 O
        • Taste bitter
        • Turn red litmus  blue
        • Act as electrolytes in solution
        • Neutralise solutions containing hydrogen ions (H + )
        • Have a slippery, ‘soapy’ feel
        • Destroy body tissue/ dissolve fatty (lipid) material
    • Measuring Acids and Bases
      • Acids and bases strengths are measured using the pH scale.
      • (pH stands for pondus hydrogenii , or potential of hydrogen.)
      • Acids have a pH between 0 and 7
      • Bases have a pH between 7 and 14
      • pH is a logarithmic scale like the Richter scale. Therefore 2 is 10 times more acidic than 3 and 10 is ten times more basic than 9.
      • 7 is pH Neutral.
    • Acid + Alkali
      • An acid will react with an alkali to form a salt and water.
      • Neutralization reactions are a special type of double replacement reactions
      • The standard reaction is:
        • acid + alkali  salt + water
      • eg:
      • hydrochloric acid + sodium hydroxide  sodium chloride + water
      • s ulfuric acid + potassium hydroxide  potassium sulfate + water
      • Write out the balanced chemical equations for each reaction
      • Write the equations for the following neutralization reactions
      • Phosphoric acid and Calcium hydroxide
      • H 3 PO 4 + Ca(OH) 2  Ca 3 (PO 4 ) 2 + H 2 O
      • Nitric acid and cesium hydroxide
      • HNO 3 + CsOH  H 2 O + CsNO 3
      Neutralization Reactions
    • Acid + Metal
      • An acid will react with a metal to form a salt and hydrogen gas.
      • The standard reaction is:
        • acid + metal  salt + hydrogen gas
      • eg:
        • hydrochloric acid + zinc  zinc chloride + hydrogen gas
        • s ulfuric acid + sodium  sodium sulfate + hydrogen gas
      • Write out the balanced chemical equations for each reaction
    • Metal + Acid
      • Reactions of metals with different acids form different salts:
        • Hydrochloric acid  chlorides
        • Sulfuric acid  sulfates
        • Nitric acid  nitrates
        • Phosphoric acid  phosphates
    • Metal Oxides
      • We have already seen how m any metals will react with oxygen to form a metal oxide:
        • eg: sodium + oxygen  sodium oxide
      • Non-metals also react with oxygen to form oxides:
        • eg: carbon + oxygen  carbon dioxide
    • Forming Basic Oxides
      • Metal oxides (that are soluble) react with water to form alkaline solutions.
      • Hence these oxides are called Basic Oxides .
        • eg: magnesium oxide + water  magnesium hydroxide
        • MgO + H 2 O  Mg (OH) 2
      • Only the very active metals form metal oxides which are soluble
    • Basic Oxides
      • Try these:
        • Sodium oxide + water 
        • Potassium oxide + water 
        • Calcium oxide + water 
      Sodium hydroxide Na 2 O + H 2 O  2NaOH Potassium hydroxide K 2 O + H 2 O  2KOH Calcium hydroxide CaO + H 2 O  Ca(OH) 2
    • Forming Acidic Oxides
      • Non-metal oxides react with water to form acidic solutions.
      • Hence these oxides are called Acidic Oxides .
        • Eg: sulfur dioxide + water  sulfurous acid
        • SO 2 + H 2 O  H 2 SO 3
    • Acidic Oxides
      • Try these:
        • sulfur tri oxide + water 
        • carbon dioxide + water 
        • nitrogen di oxide + water 
      sulfuric acid carbonic acid Nitric acid SO 3 + H 2 O  H 2 SO 4 CO 2 + H 2 O  H 2 CO 3 NO 2 + H 2 O  HNO 3
    • Acid Rain
      • Many of these non-metal oxides are being released into the atmosphere.
      • Here they are reacting with atmospheric water to form acids.
      • They then fall as Acid Rain :
        • sulfur oxides, nitrogen oxides, carbon oxides
    • Metal Oxides + Acid
      • Metal oxides will dissolve in acid to form a salt and water.
        • Eg: metal oxide + acid  salt + water
      copper oxide + sulfuric acid Copper sulfate + water CuO + H 2 SO 4 CuSO 4 + H2O
    • Metal Oxides in Acid
      • Try these:
        • zinc oxide + hydrochloric acid 
        • aluminium oxide + sulfuric acid 
        • lead oxide + sulfuric acid 
      zinc chloride + water ZnO + 2HCl  ZnCl 2 + H 2 O Al 2 O 3 + H 2 SO 4  Al 2 SO 4 + H 2 0 aluminium sulfate + water lead sulfate + water PbO + H 2 SO4  PbSO 4 + H 2 O
    • Acid + Carbonate
      • An acid will react with a carbonate to form a salt and carbon dioxide gas and water.
      • The standard reaction is:
        • acid + carbonate  salt + carbon dioxide + water
      • eg:
      • nitric acid + calcium carbonate  calcium nitrate + carbon dioxide + water
      • sulfuric acid + lead carbonate  lead sulfate + carbon dioxide + water
      • Write out the balanced chemical equations for each reaction
    • Ions in Solution
      • Acidic solutions – contain more H + than OH -
      • Basic solutions – contain more OH - than H +
      • Neutral solutions – contain equal amounts of H + and OH -
    • Autoionization of Water
      • H 2 O + H 2 O  H 3 O + + OH -
      • Water is the usual solvent for acids and bases
      • It produces equal numbers of H 3 O + and OH -
    • Redox Yr 10 Chemistry
    • Redox
      • Are reactions that involve the transfer of electrons. One species is oxidised the other is reduced
      • Oxidation – loss of electrons/addition of oxygen/ increase of ox no.
      • Reduction – gain of electrons/ loss of oxygen/ decrease of ox no.
    • Oxidisers
      • The following are oxidisers and the product that usually results
      • Oxygen water
      • Fluorine fluoride ion
      • Hydrogen ion hydrogen
      • Hydrogen peroxide water
    • Reducers
      • The following are reducers and the product that usually results
      • carbon carbon dioxide
      • Zinc Zinc ions
      • Hydrogen water
      • Sulfur dioxide sulfate ions
    • Metal + Water is redox
      • Only very active metals react with water to form an alkali and hydrogen gas.
      • The standard reaction is:
        • Metal + water  metal alkali + hydrogen gas
      • eg:
        • Sodium + water  sodium hydroxide + hydrogen gas
        • Caesium + water  caesium hydroxide + hydrogen gas
      oxidation reduction
    • Metal + Water Caesium in water
    • Complete the equation and Identify the oxidiser and reducer
      • Calcium oxide + hydrogen
      • Fluorine + hydrogen
      • Sodium hydroxide + carbon
      • Carbon monoxide + oxygen
      • Chlorine + magnesium
      • Aluminium + oxygen
    • Oxidation
      • Is the Loss of Electrons
      • is Oxidation
      • LEO
      • eg
      • Aluminium aluminium ions
    • Reduction
      • Gain of Electrons
      • Is Reduction
      • GER
      • eg
      • Iron ions iron
    • Redox
      • If both reduction and oxidation take place we can remember
            • O xidation
            • I s
            • L oss of electrons
            • R eduction
            • I s
            • G ain in electrons