C4 - Chemical changes Flashcards
How are Metal Oxides formed?
WHat are redox reactions(Oxidation and Reduction)?
- Metals react with oxygen in the air to produce metal oxides
- Oxidation and reduction involve the addition or removal of oxygen from a substance and are called redox reactions
- Oxidation is a reaction in which:
- Oxygen is added to an element or a compound
- Reduction is a reaction in which:
- Oxygen is removed from an element or a compound
- A common example is the reaction with red-brown copper metal to produce black copper oxide:
2Cu + O2 ⟶ 2CuO
- In this reaction copper metal has been oxidised since oxygen has been added to it
- Another example is the reaction of zinc oxide with carbon:
ZnO + C ⟶ Zn + CO
- In this reaction the zinc oxide has been reduced since it has lost oxygen. The carbon atom has been oxidised since it has gained oxygen
What is The reactivity series?
- The tendency of a metal to lose electrons is a measure of how reactive the metal is
- A metal that is high up on the series loses electrons easily and is thus more reactive than one which is lower down on the series
- Note that although carbon and hydrogen are nonmetals, they are included in the series as they are useful in extracting metals from their oxides by reduction processes
What happens when Metals react with Water?
metal + water → metal hydroxide + hydrogen
- Some metals react with water
- Metals above hydrogen in the reactivity series will react with water, but the reaction may be very slow
- Metals that react with cold water form a metal hydroxide and hydrogen gas:
For example calcium:
Ca + 2H2O → Ca(OH)2 + H2
calcium + water → calcium hydroxide + hydrogen
Name the Reactivity series
- There are several reactivity series mnemonics to help you remember the order of the metals
- One that we like goes as follows: “
Please
send
lions,
cats,
monkeys
and
cute
zebras
into
hot
countries
signed
Gordon”
How do Metals react with Acids?
- Only the metals below hydrogen in the reactivity series will not react with acids
- When acids and metals react, the hydrogen atom in the acid is replaced by the metal atom to produce a salt and hydrogen gas:
metal + acid → metal salt + hydrogen
- For example iron:
Fe + 2HCI → FeCl2 + H2
iron + hydrochloric acid → iron(II)chloride + hydrogen
- In both these types of reactions (water and acids) the metals are becoming positive ions
- The reactivity of the metals is related to their tendency to become an ion
- The more reactive the metal the more easily it becomes an ion (by losing electrons)
Why do Non-metals appear in the Reactivity Series?
- Why do non-metals appear in the reactivity series of metals?
- A reactivity series will usually contain the elements carbon and hydrogen
- This is beause these elements play different roles in our understanding the reactions of metals and our ability to predict how metals can be extracted from their ores
- From the reactions with water and acids we have seen that whether a reaction takes place depends on the position of the metal in the reactivity series relative to hydrogen
- A reaction takes place if the metal is able to displace hydrogen from water or acids
-
Carbon is a cheap reducing agent which can be used to remove oxygen from metal oxide ores
- Placing carbon in the reactivity series allows us to see whether a metal oxide can be reduced or not by carbon
- Metals below carbon can be extracted by heating the oxide with carbon
- Metals higher than carbon have to be extracted by other methods, such as electrolysis
What is a Displacement reaction?
- The reactivity of metals decreases going down the reactivity series.
- This means that a more reactive metal will displace a less reactive metal from its compounds
- Two examples are:
- Reacting a metal with a metal oxide (by heating)
- Reacting a metal with an aqueous solution of a metal compound
- For example it is possible to reduce copper(II) oxide by heating it with zinc.
- The reducing agent in the reaction is zinc:
Zn + CuO → ZnO + Cu
zinc + copper oxide → zinc oxide + copper
What happens in displacement reactions between metals & Aqueous soloutions?
- This is easily seen as the more reactive metal slowly disappears from the solution, displacing the less reactive metal
- For example, magnesium is a reactive metal and can displace copper from a copper sulfate solution:
Mg + CuSO4→ MgSO4 + Cu
- The blue color of the CuSO4 solution fades as colorless magnesium sulfate solution is formed.
- Copper coats the surface of the magnesium and also forms solid metal which falls to the bottom of the beaker
What are Redox reactions?
balanced equation for the reaction between magnesium and copper sulfate solution can be written in terms of the ions involved:
Mg(s) + Cu2+(aq) + SO42-(aq) → Mg2+(aq) + SO42-(aq) + Cu(s)
Sulfate ions, SO42-, appear on both sides of the equation, but they do not take part in the reaction. The equation can be rewritten without them:
Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)
This equation is an example of a balanced ionic equation. It can be split into two half equations :
Mg(s) → Mg2+(aq) + 2e- (oxidation)
Cu2+(aq) + 2e- → Cu(s) (reduction)
Notice that:
- magnesium atoms lose electrons - they are oxidised
- copper ions gain electrons - they are reduced
Reduction and oxidation happen at the same time, so the reactions are called redox reactions.
Oxidation is the ______ of Electrons, Reduction is _______ of electrons
Oxidation is the loss of electrons, and reduction is the gain of electrons.
It helps to remember OIL RIG - oxidation is loss of electrons, and reduction is gain of electrons.
Displacement reactions are just one example of redox reactions. Electrolysis reactions are also redox reactions.
How do we extract Metals?
- The Earth’s crust contains metals and metal compounds such as gold, copper, iron oxide and aluminium oxide
- Useful metals are often chemically combined with other substances forming ores
- A metal ore is a rock that contains enough of the metal to make it worthwhile extracting
- They have to be extracted from their ores through processes such as electrolysis, using a blast furnace or by reacting with more reactive material
- In many cases the ore is an oxide of the metal, therefore the extraction of these metals is a reduction process since oxygen is being removed
- Common examples of oxide ores are iron and aluminium ores which are called haematite and bauxite respectively
- Unreactive metals do not have to be extracted chemically as they are often found as the uncombined element
- This occurs as they do not easily react with other substances due to their chemical stability
- Examples include gold and platinum which can both be mined directly from the Earth’s crust
How are Elements more reactive then Carbon extracted?
How are elements less reactive than Carbon but more reactive than silver Extracted?
How are Elements less copper extracted?
- The most reactive metals are at the top of the series
- The tendency to become oxidised is thus linked to how reactive a metal is and therefore its position on the reactivity series
- Metals higher up are therefore less resistant to oxidation than the metals placed lower down which are more resistant to oxidation
- The position of the metal on the reactivity series determines the method of extraction
- Higher placed metals (above carbon) have to be extracted using electrolysis as they are too reactive and cannot be reduced by carbon
- Lower placed metals can be extracted by heating with carbon which reduces them
How is Iron Extracted?
Extracting iron
Iron(III) oxide is reduced to molten iron when it reacts with carbon. One of the products is carbon monoxide:
iron(III) oxide + carbon → iron + carbon monoxide
Fe2O3(s) + 3C(s) → 2Fe(l) + 3CO(g)
This method of extraction works because carbon is more reactive than iron, so it can displace iron from iron compounds. Extracting a metal by heating with carbon is cheaper than using electrolysis.
How is Aluminium Extracted?
Extracting aluminium
Aluminium is more reactive than carbon so it must be extracted from its compounds using electrolysis. Even though aluminium is more abundant than iron in the Earth’s crust, aluminium is more expensive than iron. This is mainly because of the large amounts of electrical energy used in the extraction process.
Electrolysis of aluminium oxideThe electrolyte
Aluminium ore is treated to produce pure aluminium oxide. The electrolytes used in electrolysis are ionic compounds:
- in the molten state, or
- dissolved in water
Aluminium oxide is insoluble in water, so it must be molten to act as an electrolyte. However, the melting point of aluminium oxide is high. A lot of energy must be transferred to break its strong ionic bonds, and this is expensive. To reduce costs, powdered aluminium oxide is dissolved in molten cryolite. This ionic compound melts at a lower temperature than aluminium oxide, reducing costs. However, significant amounts of energy are required to melt the cryolite.
The electrolysis process
The diagram shows an electrolysis cell used to extract aluminium. Both electrodes are made of graphite, a form of carbon with a high melting point and which conducts electricity.
During electrolysis:
- at the cathode, aluminium ions gain electrons and form aluminium atoms
- at the anode, oxide ions lose electrons and form oxygen gas
The oxygen reacts with the carbon anodes, forming carbon dioxide. So the anodes are gradually oxidised. They must be replaced frequently, adding to the cost of producing aluminium.
Worked example - Higher
Explain, with the help of a half equation, how oxide ions are oxidised during the electrolysis of aluminium oxide.
The half equation is: 2O2- → O2 + 4e-. It shows that oxide ions lose electrons, and oxidation is loss of electrons.
What are Acids?
Acids
Acids form acidic solutions in water. Acids produce hydrogen ions, H+ in aqueous solution. For example:
HCl(aq) → H+(aq) + Cl-(aq)
Acidic solutions have pH values less than 7.
What are Alkalis?
Alkalis form alkaline solutions in water. Alkalis produce hydroxide ions, OH- in aqueous solution. For example:
NaOH(aq) → Na+(aq) + OH-(aq)
Alkaline solutions have pH values greater than 7.
What are Neutral Soloutions?
A neutral solution is neither acidic, nor alkaline. A neutral solution has a pH value of 7.
What is the pH scale?
The pH scale measures the acidity or alkalinity of a solution. The pH of a solution can be measured using a pH probe, or estimated using universal indicator and a colour chart.
Universal indicator is one example of an acid-alkali indicator. Indicators show whether a solution is acidic, neutral (pH 7) or alkaline. The table shows the colours for litmus paper.
What happens when Acids with Metals?
- Only metals above hydrogen in the reactivity series will react with dilute acids.
- The more reactive the metal then the more vigorous the reaction will be.
- Metals that are placed high on the reactivity series such as potassium and sodium are very dangerous and react explosively with acids.
- When acids react with metals they form a salt and hydrogen gas:
- The general equation is:
metal + acid ⟶ salt + hydrogen
- Some examples of metal-acid reactions and their equations are given below:
What happens when Metal Hydroxides and Metal oxides react with an acid?
Reactions of Acids with Metal Oxides and Metal Hydroxides
- Metal oxides and metal hydroxides act as bases
- When they react with acid, a neutralisation reaction occurs
- In all acid-base neutralisation reactions, salt and water are produced
- The following are some specific examples of reactions between acids and metal oxides / hydroxides:
2HCl + CuO ⟶ CuCl2 + H2O
H2SO4 + 2NaOH ⟶ Na2SO4 + 2H2O
HNO3 + KOH ⟶ KNO3 + H2O
What happens when Acids react with Metal Carbonates?
- Acids will react with metal carbonates to form the corresponding metal salt, carbon dioxide and water
- These reactions are easily distinguishable from acid – metal oxide/hydroxide reactions due to the presence of effervescence caused by the carbon dioxide gas
The following are some specific examples of reactions between acids and metal carbonates:
2HCl + Na2CO3 ⟶ 2NaCl + H2O + CO2
Reactions with carbonates
A salt, water and carbon dioxide are produced when acids react with carbonates. In general:
Acid + carbonate → salt + water + carbon dioxide
For example:
Hydrochloric acid + copper carbonate → copper chloride + water + carbon dioxide
2HCl(aq) + CuCO3(s) → CuCl2(aq) + H2O(l) + CO2(g)
H2SO4 + CaCO3⟶ CaSO4 + H2O + CO2
How do you predict the names of the salt produced?
Salts
- A salt is a compound that is formed when the hydrogen atom in an acid is replaced by a metal
- For example if we replace the H in HCl with a potassium atom, then the salt potassium chloride is formed, KCl
- Salts are an important branch of chemistry due to the varied and important uses of this class of compounds
- These uses include fertilisers, batteries, cleaning products, healthcare products and fungicides
Naming salts
- The name of a salt has two parts
- The first part comes from the metal, metal oxide or metal carbonate used in the reaction
- The second part comes from the acid
- The name of the salt can be determined by looking at the reactants
- For example hydrochloric acid always produces salts that end in chloride and contain the chloride ion, Cl–
- Other examples:
- Sodium hydroxide reacts with hydrochloric acid to produce sodium chloride.
- Zinc oxide reacts with sulfuric acid to produce zinc sulfate
- A list of the common ions and their formulae is shown below
Salts have no overall charge since the sum of the charges on the ions is equal to zero
- From the table of the common ions it is clear the charges on each the Group I elements is always +1, Group II is +2, Group VI is -2 and Group VII is -1
- If you know the ions present in a salt you can identify the formula from balancing the charges
What is the difference between a Strong and Weak Acid?
Strong and weak acids
Acids in solution are a source of hydrogen ions, H+. The hydrogen ions are produced when the acid dissociates or breaks down to form ions.
Strong acids
Strong acids completely dissociate into ions in solution. For example, hydrochloric acid is a strong acid. It ionises completely to form hydrogen ions and chloride ions:
HCl(aq) → H+(aq) + Cl-(aq)
Nitric acid and sulfuric acid are also strong acids.
Weak acids
Weak acids only partially dissociate in solution. For example, ethanoic acid is a weak acid. It is only partially ionised to form hydrogen ions and ethanoate ions:
CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)
The ⇌ symbol is used in the equation to show that the reaction is a reversible reaction and does not go to completion.
What is a Titration?
- Titrations are a method of analysing the concentration of solutions
- Acid-base titrations are one of the most important kinds of titrations
- They can determine exactly how much alkali is needed to neutralise a quantity of acid – and vice versa
- You may be asked to calculate the moles present in a given amount, the concentration or volume required to neutralise an acid or a base
- Titrations can also be used to prepare salts or other precipitates and in redox reactions
- Indicators are used to show the endpoint in a titration
- Wide range indicators such as litmus are not suitable for a titration as they do not give a sharp enough colour change at the end point
- Some of the most common indicators with their corresponding colours are shown below
What do you have to calculate in Titrations?
- Once a titration is completed and the average titre has been calculated, you can now proceed to calculate the unknown variable using the formula triangle as shown below
How do you make salts from Acids and Alkalis?
A soluble salt can be prepared by reacting an acid with a dilute solution of an alkali such as sodium hydroxide or ammonia. The main steps are:
- Carry out a titration. This is to determine the volumes of acid and alkali that must be mixed to obtain a solution containing only salt and water.
- Mix the acid and alkali in the correct proportions, as determined in step 1, but this time without including an indicator.
Pure dry crystals can be produced by crystallisation, followed by drying on a watch glass or in a warm oven.
Carrying out a titration to find out volumes of acid and alkali solutions that reactApparatus
The apparatus needed includes:
- a pipette to accurately measure the volume of a reactant before transferring it to a conical flask
- a burette to add small, measured volumes of one reactant to the other reactant
Method
This is an outline method for carrying out a titration in which an acid is added to an alkali. The method is the same for sulfuric acid, hydrochloric acid and nitric acid.
- Use the pipette and pipette filler to add a measured volume of sodium hydroxide solution to a clean conical flask.
- Add a few drops of indicator and put the conical flask on a white tile.
- Fill the burette with hydrochloric 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 (when the indicator first permanently changes colour). Note the final volume reading.
- Repeat steps 1 to 5 until concordant titres are obtained. More accurate results are obtained if acid is added drop by drop near to the end-point.
Readings should be recorded to two decimal places, ending in 0 or 5 (where the liquid level is between two graduations on the burette). The titre is the volume added (the difference between the end and start readings).
Analysis
At least two concordant titres should be ticked (✔) in the table above. These are titres within 0.20 cm3 (or sometimes 0.10 cm3) of each other.
Worked example
Calculate the mean titre from the table above. Ignore the rough run, and run 2 (because they are not concordant):
Mean titre = (24.60 + 24.70)/2
= 24.65 cm3
What are Electrolytes?
What is an Electrode?
What is an Anode?
What is an Anion?
What is a cathode?
What is a Cation?
- Covalent compounds cannot conduct electricity hence they do not undergo electrolysis
- An electrolytic cell is the name given to the set-up used in electrolysis and which consists of the following:
- Electrode: a rod of metal or graphite through which an electric current flows into or out of an electrolyte
- Electrolyte: ionic compound in molten or dissolved solution that conducts the electricity
- Anode: the positive electrode of an electrolysis cell
- Anion: negatively charged ion which is attracted to the anode
- Cathode: the negative electrode of an electrolysis cell
- Cation: positively charged ion which is attracted to the cathode
Where do the ions move in Electrolysis?
- During electrolysis the electrons move from the power supply towards the cathode
- Electron flow in electrochemistry thus occurs in alphabetical order as electrons flow from the anode to the cathode
- Positive ions within the electrolyte migrate towards the negatively charged electrode which is the cathode
- Negative ions within the electrolyte migrate towards the positively charged electrode which is the anode
How are the products of electrolysis formed?
Products of electrolysis
When ions reach an electrode, they gain or lose electrons. As a result, they form atoms or molecules of elements:
- positive ions gain electrons from the negatively charged cathode
- negative ions lose electrons at the positively charged anode
Molten lead bromide, PbBr2(l), is an electrolyte. During electrolysis:
- Pb2+ ions gain electrons at the cathode and become Pb atoms
- Br- ions lose electrons at the anode and become Br atoms, which pair up to form Br2 molecules
So lead forms at the negative electrode and bromine forms at the positive electrode.
Example
Predict the products of electrolysis of molten calcium chloride.
Positively charged calcium ions move to the negative electrode. Here, they gain electrons to form calcium atoms, so calcium is formed at the negative electrode.
Negatively charged chloride ions move to the positive electrode. Here, they lose electrons to form chlorine atoms. The atoms join up in pairs to form Cl2 molecules, so chlorine gas is formed at the positive electrode.
During the electrolysis of molten salts, a metal forms at the cathode and a non-metal forms at the anode.
How do you use electrolysis on Lead Bromide?
- Lead(II) bromide is a binary ionic compound meaning that it is a compound consisting of just two elements joined together by ionic bonding
- When these compounds are heated beyond their melting point, they become molten and can conduct electricity as their ions can move freely and carry the charge
- These compounds undergo electrolysis and always produce their corresponding element
- To predict the products of any binary molten compound first identify the ions present
- The positive ion will migrate towards the cathode and the negative ion will migrate towards the anode
- Therefore the cathode product will always be the metal and the product formed at the anode will always be the non-metal
Method:
- Add lead(II) bromide into a crucible and heat so it will turn molten, allowing ions to be free to move and conduct an electric charge
- Add two graphite rods as the electrodes and connect this to a power pack or battery
- Turn on the power pack or battery and allow electrolysis to take place
- Negative bromide ions move to the positive electrode (anode) and lose two electrons to form bromine molecules. There is bubbling at the anode as brown bromine gas is given off
- Positive lead ions move to the negative electrode (cathode) and gain electrons to form grey lead metal which deposits on the bottom of the electrode
Electrode Products:
Anode: Bromine gas
Cathode: Lead metal
How do determine the Method of Extraction?
- The position of the metal on the reactivity series determines the method of extraction
- Higher placed metals (above carbon) have to be extracted using electrolysis as they are too reactive and cannot be reduced by carbon
- Lower placed metals can be extracted by heating with carbon which reduces
- Electrolysis is very expensive as large amounts of energy are required to melt the ores and produce the electrical current
- The reactivity series of metals is shown below with the corresponding method of extraction
How is Aluminium Extracted?
Raw Materials:
- Aluminium Ore (bauxite)
- Cryolite (sodium aluminium fluoride)
Explanation:
- The bauxite is first purified to produce aluminium oxide, Al2O3
- Aluminium oxide has a very high melting point so it is first dissolved in molten cryolite producing an electrolyte with a lower melting point, as well as a better conductor of electricity than molten aluminium oxide
- This reduces the costs considerably making the process more efficient
- The electrolyte is a solution of aluminium oxide in molten cryolite at a temperature of about 1000 °C
- The molten aluminium is siphoned off from time to time and fresh aluminium oxide is added to the cell
- The cell operates at 5-6 volts and with a current of 100,000 amps
- The heat generated by the huge current keeps the electrolyte molten
- A lot of electricity is required for this process of extraction which is a major expense
- The overall equation is:
2Al2O3 (l) ⟶ 4Al (l) + 3O2 (g)
- Some of the oxygen produced at the positive electrode then reacts with graphite (carbon) electrode to produce carbon dioxide gas:
C (s) + O2 (g) ⟶ CO2 (g)
- This causes the carbon anodes to burn away, so they must be replaced regularly
