(2d) Oxygen and Oxides

Oxygen and Oxides - Section 2d 
2.16 recall the gases present in air and their approximate percentage by volume 

Nitrogen = 78% 
Oxygen = 21%
Argon = nearly 1%
Co2 (carbon dioxide) = o.o4% 

2.17 explain how experiments involving the reactions of copper, iron and phosphorous with air can be used to investigate the percentage of volume of oxygen in the air 

The proportion of oxygen in the atmosphere can be measured through two separate experiments, one using copper and one using iron and phosphorus. 

Using copper 

• When heated, copper reacts with oxygen in the air to make copper (II) oxide. So this reaction uses up oxygen from the atmosphere. 
• Pass a known volume of air (say 100cm^3) over hot copper in a tube using two syringes. The markers on the syringes will tell you how much oxygen has been used up. 
• If you started with 100cm^3 of air, and after passing it over the hot copper you see that 20cm^3 of air has gone, then you know that that 20cm^3 would have been oxygen (oxygen is used up when copper reacts to make copper.) 
• So, you could say that about one fifth of the air (20%) is oxygen. 

Using Iron or Phosphorus

• Iron reacts with oxygen in the air to form rust.
• Soak some iron wool in acetic acid, which will catalyse the reaction. Push the wool into the test tube, covering the of the tube and then inverting it into a beaker of water. 
• Over time, the level of water in the tube will rise. This is because the iron reacts with the oxygen in the air, making iron oxide. The water rises to fill the space that the oxygen took up. 
• In order to work out the percentage of the air that is oxygen, you need to make the starting and finishing position of the water. 
• Fill the tube up to each mark with water and pour the contents into a measuring cylinder to find the volume of air at the start and at the end. 
• The difference between the volumes at the start and at the end will give you the percentage of the initial volume that has been used up - this should be around 21%. This tells you how much oxygen was present in the air as it is the oxygen that reacted with the rust and therefore decreased the volume. 

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese (IV) oxide as a catalyst 

You can make pure oxygen in the lab from hydrogen peroxide (H2O2).

Hydrogen peroxide is put in a conical flask with manganese (IV) oxide as a catalyst. The hydrogen peroxide will decompose into water and oxygen. Normally, this reaction would be slow but it can be sped up with the presence of manganese (IV) oxide, which will speed the reaction up without being used up in it. Then, you can collect the oxygen that is produced over water (using a delivery tube to bubble the gas into an upside-down measuring cylinder or gas jar filled with water,) or simple use a gas syringe to collect it. 

Equation for decomposition of hydrogen peroxide into oxygen and water;

word equation:                       hydrogen peroxide > water + oxygen 

balanced symbol equation:            2H2O2 (aq)      > 2HO (l)  + O2 (g) 

2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced

When an element is burnt in air it reacts with the oxygen to form an oxide. The oxides will either have an acidic character, or a basic character. 

Magnesium 

• Burns with a bright white flame in air. 
• White powder is formed, this is magnesium oxide. 
• Magnesium oxide is slightly alkaline when dissolved in water, so it has a basic character. 
• Equation: 2Mg (s) + O2 (g) > 2MgO (s)

Bright white flame that magnesium burns with when burnt in air.

Carbon 

• Will burn in  air if it is very strongly heated. 
• Burns with an orangey/yellow flame. 
• Produces carbon dioxide gas. 
• Carbon dioxide is slightly acidic when dissolved in water, it has an acidic character. 
• Equation: C (s) + O2 (g)  > CO2 (g)

Carbon burning in air with a yellow/orange flame. 

Sulfur

• Burns in air with a pale blue flame. 
• Produces sulfur dioxide. 
• Sulfur dioxide is slightly acidic when dissolved in water, it has an acidic character. 
• Equation: S (s) + O2 (g) > SO2 (g)

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

• Marble chips (a metal carbonate) are placed at the bottom of a flask, and dilute hydrochloric acid is added. 
• The dilute hydrochloric acid (HCI) reacts with the calcium carbonate to produce calcium chloride, water and carbon dioxide gas. 
• The carbon dioxide gas is the collected in a gas syringe or using downward delivery. 
• Equation: hydrochloric acid + calcium carbonate > calcium chloride + water + carbon dioxide.

2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper (II) carbonates. 

Another way of making CO2 is by heating a metal carbonate, a process known as thermal decomposition (when a substance breaks down into simpler substances when heated.) Copper (II) carbonate is a green powder that will easily decompose upon being heated, forming carbon dioxide and copper (II) oxide. Equation for this reaction is as follows:

• CuCO3 (s) > CuO (s) + CO2 (g)

c) Group 7 elements - chlorine, bromine and iodine

c) Group 7 elements - chlorine, bromine and iodine 

2.8 recall the colours and physical states of the elements at room temperature. 

Group 7 Element	Colour	Physical state (at room temperature)
Chlorine	Green	Gas
Bromine	Red/brown	Liquid
Iodine	Dark grey	Solid

2.9 make predictions about the properties of other halogens in this group

As you travel down Group 7, the reactivity of the elements decreases. The elements will therefore have a darker colour and higher boiling point as you travel down the group, so fluorine will be the most reactive as it is the highest element in the group. It will have the darkest colour and the highest  boiling point. Then, the element furthest down in the group, Astatine, will have the lightest colour and the lowest  boiling point as it has the largest atomic number of all the elements in the group (the shell with the missing electron is furthest from the nucleus, so the pull from the positive nucleus is weaker.) 

2.10 understand the difference between hydrogen chloride gas and hydrochloric acid 

 Hydrogen chloride (HCI) is a gas at room temperature. Hydrochloric acid is hydrogen chloride dissolved in water. When dissolved in water, the HCI molecules split up into H+ ions and CI- ions in a process known as dissociation. The solution that is formed as a result of this is hydrochloric acid. Therefore, hydrogen chloride dissociates in water to form a solution of hydrochloric acid (it is an acid because it has H+ ions.) 

2.11 explain, in terms of disassociation, why hydrogen chloride is acidic in water but not in methyl benzene

Hydrogen chloride is acidic in water because it's molecules split into H+ and CI- ions in water, making it an acidic solution due to the presence of H+ ions. However, hydrogen chloride is not acidic in methyl benzene because it does not disassociate into H+ and CI- ions. With no H+ ions present, the hydrogen chloride is not an acid in methyl benzene. 

2.12 describe the relative reactivites of the elements in group 7

Group 7 elements become less reactive as you go down the group. The higher up Group 7 an element is, the closer the shell with the missing electron is to the nucleus, therefore the stronger the pull from the positive nucleus. The further down the element in Group 7, the further the electron shell with the missing electron and therefore the weaker the pull from the nucleus. 

2.13 describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

A displacement reaction is one in which a more reactive element displaces a less reactive element from a compound. The elements in group 7 take part in these, following the rule that the elements are more reactive the higher up in the group they are. A more reactive halogen will displace a less reactive one that is bonded as a salt, but only if the salt is dissolved in water or a gas.

So for example chlorine is more reactive than iodine as it is higher up in Group 7. Therefore, if you add chlorine water to potassium iodide the more reactive chlorine will react with the potassium in the potassium chloride to displace the iodide, forming potassium chlorine. 

2.14 understand these displacement reactions as redox reactions

A redox reaction is a reaction in which both reduction (gain of electrons) and oxidation (loss of electrons) happens simultaneously. Displacement reactions between halogens and salt solutions are redox reactions because both oxidation and reduction occurs. 

So, in the reaction between chlorine water and potassium iodide, chlorine is reduced because it gains electrons and iodine is oxidised because it loses electrons. The displacement is therefore a redox reaction. 

2b) Group 1 elements - lithium, sodium and potassium

2b) Group 1 elements - lithium, sodium and potassium 

2.6 describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements 

Group 1 elements (the alkali metals) all have one electron in their outer shell. As they are in the same group, they have similar chemical properties and will react with substances (including water) in a similar way. The Group 1 elements of lithium, sodium and potassium all react in a similar way with water, they react very vigorously which indicates that they are a family of elements and must have a similar electronic configuration. Although all the alkali metals will react similarly with water, as you go down Group 1 the elements become more reactive and will react more vigorously with water. 

Lithium: 

• Lump of lithium in water will move slowly around the surface, fizzing until it disappears. The water will become alkaline and the indicator solution will turn purple because of this. It will take approximately 30 seconds for this reaction to occur. Lithium is at the top of the group and is therefore the least reactive of all the alkali metals.

Lithium reacting with water. It will fizz like this until it disappears. 

Sodium; 

• The lump of sodium will fizz rapidly in the water and will move quickly around the surface.The indicator solution will turn purple as the water will become alkaline and it may even ignite. The time for this reaction to occur will be approximately 20 seconds. Sodium is further down in Group 1 than lithium, hence why it reacts more vigorously with the water and why it takes less time for this to happen. 

Sodium reacting with water. Having fizzed and moved quickly around the surface of the water,it has now ignited.

Potassium: 

• Lump of potassium will react vigorously with the water, and will burn with a lilac flame, or even explode. The indicator solution will again turn purple as the water has become alkaline. This reaction will occur in approximately 5 seconds. Potassium is the furthest down in Group 1 of all the three elements mentioned, hence why it reacts most vigorously with water and also why it has the quickest reaction time. 

                                                          

Potassium reacting with water. 

2.7 describe the relative reactivites of the elements in group 1

Elements in group 1 become more reactive as you go down the group. So, the least reactive metal in Group 1 is also the highest element in the group,so it would be lithium, then sodium, then potassium and so on. 

2.8 explain the relative reactivites of the elements in Group 1 in terms of distance between the outer electrons and the nucleus. 

The further you go down Group 1, the higher the period number of each element is. While the group number stays the same, the period number will increase. This means that as you go down the group, the elements have a higher period number, meaning they have more energy levels (shells.) Because of this, the elements' outer electrons will be further away from the nucleus as you go down the group, so the element will be able to lose electrons far easier as the attraction between the outermost and the nucleus becomes less. Therefore, the outer electron will be more easily lost and the element will be more reactive as you travel down Group 1.

2a) The Periodic Table

Section 2: Chemistry of the elements a) The Periodic Table 

2.1 understand the terms group and period 

Group = 

• The columns of the periodic table.
•  The number of the group an element is in tells you how many electrons it has in its outer shell (eg. If an electron is in group 3, it has 3 electrons in its outer shell.)
• Elements in the same group have similar chemical properties (same amount of electrons in their outer shell, they will therefore react and bond in similar ways)
• Properties of elements change as you go down a group. 

Period = 

• The rows of the periodic table. 
• The number of the period an element is in tells you how many energy levels (shells) it has. For example, if an element is in period 5 then it has 5 energy levels/ shells.
• Properties of elements change as you go along a period. 

2.2 recall the positions of metals and non-metals in the periodic table 

The elements on the left of the zig-zag are all metals. 

The elements on the right of the zig-zag are all non-metals.

2.3 explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides 

Metals: 

• Metals are good conductors of electricity because they allow charge to pass through them easily. (They have delocalised electrons which make it possible for an electrical current to pass through.) 
• Metal oxides are basic, and will therefore neutralise acids. Metal oxides which dissolve will form solutions with a pH higher than 7 (alkalis).

Non-metals 

• Non-metals are poor conductors of electricity because they are held tightly together by very strong covalent bonds, and are not free to move around or pass an electric current. 
• Non-metal oxides are acidic. They dissolve in water to form solutions with a pH of less than 7 (acidic.)

2.4 understand why elements in the same group of the Periodic Table have similar chemical properties 

Elements in the same group of the Periodic Table all have the same number of electrons in their outer shell. (Group number = number of electrons in outer shell.) Therefore, all of the elements in a group will have the same amount of electrons they need to gain/lose in order to become stable, meaning they will all react and bond with other elements very similarly and so have similar chemical properties. 

2.5  understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations. 

All of the noble gases have full outer shells of electrons, so they are stable.  Because they are stable, they do not need to react with anything or form a bond, so they are also inert (unreactive). For example, helium is a noble gas that has 2 electrons in its outer shell (it only has one shell) , and is its therefore fully stable and inert. 

i) Electrolysis

i) Electrolysis

1.48 understand that an electric current is a flow of electrons or ions
An electric current is a flow of electrons. In order for a substance to conduct electricity, it must have charged particles that are free to move around, and as electrons and ions both fit this description, they can conduct electricity. 

1.49 understand why covalent compounds do not conduct electricity
Covalent compounds do not contain charged particles as they share electrons instead of losing or gaining them. Therefore, they do not have the mobile charged particles they need in order for an electric current to be present, so they do not conduct electricity. 

1.50 understand why ionic compounds conduct electricity only when molten or in

solution
Ionic compounds do not conduct electricity when solid as they are not free to move around. However, they do conduct electricity when molten or in solution as the ions are free to move around and able to move to the oppositely charged electrode. 

1.51 describe experiments to distinguish between electrolytes and non electrolytes
1)Set up an electric circuit with an LED and a break in the wire.
2) Put both ends of wire into a solution/molten substance.  

Then, if the LED lights up there is a current flowing as this will only occur if the solution is conducting, so it must be an electrolyte. On the other hand, if the LED does not light up then this means that there is no current flowing and the solution is not conducting electricity, so it must be a non electrolyte.

1.52 understand that electrolysis involves the formation of new substances when
ionic compounds conduct electricity

In order for electrolysis to occur, ionic compounds must conduct electricity. The positively charged ions (cations) will move to the negative electrode (cathode) and the negatively charged ions (anions) will move to the positive electrode (anode.) Having moved to their respective electrodes, the ions will then become atoms as electrons will be lost at the cathode and gained at the anode, reverting them back to molecules and therefore forming new substances. 

1.54 describe experiments to investigate electrolysis, using inert electrodes, of aqueous solutions such as sodium chloride, copper (II) sulfate and dilute sulfuric acid and predict the products.

• An inert electrode is one that helps the electrolysis but is not used up in the reaction itself. 
• In aqueous solutions, there are H+ (hydrogen) and OH- (hydroxide) ions present from the water, as well as the ions from the ionic compound. 
• At the cathode, if H+ ions and metal ions are present, hydrogen gas will be produced if the metal ions are more reactive than the H+ ions. If the H+ ions are more reactive than the metal ions, then a solid layer of the pure metal will be produced instead.
• At the anode, if OH- and halide ions (Cl-, Br-, I-) are present, molecules of chlorine, bromine or iodine will be formed. If there are no halide ions present, then oxygen will be formed. 

For example, a solution of sodium chloride (NaCl) contains four different ions: Na+ and Cl- from the ionic compound and OH- and H+ ions from the water. 

• At the cathode, negatively charged ions lose electrons. These negatively charged ions are the sodium ions, which lose electrons that are given to the hydrogen ions. Therefore, the sodium (metal) is the more reactive of the two and following the rule above, this means that we can expect hydrogen gas to be produced. 
• At the anode, positively charged electrons gain ions. These positively charged ions are the hydroxide ions, which gain electrons as the chloride ions lose them. As there are both OH- ions and halide ions (Cl- in this case) present, then we can expect chlorine gas to be produced at the anode. 

1.55 write ionic half-equations representing the reactions at the electrodes during electrolysis

A ionic half-equation shows you what happens at one of the electrodes during electrolysis. A half-equation should be balanced by adding or taking away the number of electrons equal to the number of charges on the ions in the equation. 

So, Pb2+ 2e- →  Pb (electrons are represented by 'e' in the equation.) 

This is the ionic half-equation that shows what happens at the cathode when molten lead bromide (PbBr2) is electrolysed. The positive Pb+ ions are attracted to the negative cathode., where a lead ion accepts two electrons to become a lead ion. This will form molten lead that will sink to the bottom.

The equation for the reaction at the anode = 2Br- → Br2 + 2e-

This negative Br- ions are attracted to the positive anode, where two bromide ions lose one electron each and become a bromine molecule. This causes brown bromide gas to form at the top of the anode. 

1.56 recall that one faraday represents one mole of electrons 

96500 coulombs is one faraday. One faraday contains one mole of electrons.

1.57 calculate the amounts of the products of the electrolysis of molten salts and aqueous solutions

Some molten lead (II) chloride (PbCl2) is electrolysed for 20 minutes. The current flowing is 5 amps. Find the mass of the lead produced. 

1) Write out the ionic half-equation: Pb2+ +2e-  → Pb 

2) Calculate the number of faradays: Coloumbs = current (amps) x time (s) 

                                                                            = 5 x (20x60)  (20 x 60 to find number of time in seconds overall)

                                                                             = 6000 coloumbs 

 Number of faradays = coloumbs / 96000 (96000 coloumbs = one faraday)

                                = 0.0625 faradays

3) Calculate the number of moles of lead produced. To do this, you must divide the number of faradays by the number of electrons in the half-equation.) 

So, faradays / number of electrons = 0.0625/2

                                                           = 0.02125 moles of lead atoms. 

4) Substitute the Mr values from the periodic table to work out the mass of solid lead produced. 

Mass of lead = Mr x No. of moles 

                      = 207 x 0.03125 

                      = 6.5g (1d.p.)

h) Metallic crystals

1.46 understand that a metal can be described as a giant structure of positive
ions surrounded by a sea of delocalised electrons

Metals have a giant structure of positive ions, surrounded by a 'sea of delocalised electrons.' So, a metal is a lattice of positive ions in a sea of delocalised, or free, electrons. These electrons are not attached to a particular atom and can move freely. 

1.47 explain the electrical conductivity and malleability of a metal in terms of its

structure and bonding.

As metals have delocalised electrons and electrons carry electricity, they are good conductors of electricity. The free electrons carry electrical current through the material and do this easily.

The structure of a metal is composed of rows of atoms on top of eachother, The atoms are all the same size and the layers will therefore be able to slide easily over one another, making them malleable. 

Covalent Substances (Section 1g)

Covalent Substances (Section 1g) 

1.38 describe the formation of a covalent bond by the sharing of a pair of 
electrons between two atoms 

A covalent bond is a bond formed between atoms by sharing electrons (one each) with other atoms. This leaves them stable, as they now have a full outer shell.

For example, Hydrogen has just one electron in their first and only shell, and need one more to become stable. Chlorine has seven electrons in it's outer shell and therefore needs only one more electron, so the two atoms form a bond in which they share an electron, which gives hydrogen and chlorine the extra electron it needs for a full outer shell. Once bonded in this manner, hydrogen and chlorine form hydrogen chloride (HCl).

 1.39 understand covalent bonding as a strong attraction between the bonding 
pair of electrons and the nuclei of the atoms involved in the bond 

There is a strong attraction between the electrons being shared by atoms in a covalent bond and the nuclei of all atoms involved in the bond. This is because the nucleus has a positive charge, (composed of protons and neutrons) and the electrons have a negative charge. 

1.40 explain, using dot and cross diagrams, the formation of covalent compounds 

by electron sharing for the following substances; i) hydrogen, ii) chlorine, iii) hydrogen chloride, iv) water, v) methane, vi) ammonia, vii) oxygen, viii) nitrogen, ix) carbon dioxide, x) ethane, xi) ethene. 

A dot and cross diagram is constructed by drawing the outer shells of the atoms that are bonding, with an overlap in which the electrons they share should lie. All of the electrons on one of the shells should be dots, and crosses on the other so it is clear to see which electrons belong to each atom. 

i) Hydrogen (H2) 

Hydrogen atoms have just one electron, and need one more to become stable. Two hydrogen atoms can form a covalent bond to become stable. They are then known as hydrogen, or H2. 

ii) Chlorine (Cl2)

Chlorine atoms each have seven electrons in their outer shell. Because of this, they each need one more electron to have a full outer shell and become stable. Two chlorine atoms can react to become stable and form Cl2. 

iii) Hydrogen chloride (HCl)

As hydrogen and chlorine both need only one electron to have a full outer shell, they can react and become stable. This forms hydrogen chloride, or HCl. 

iv) Water (H2O)

Oxygen has six electrons in it's outer shell and needs two more to become full. An oxygen atom can react with two hydrogen atoms (both need one more electron) to become stable. It has now become water (H2O).


v) Methane (CH4)

Carbon has four outer electrons, so needs four more to become full. Hydrogen atoms each have one electron, so four of these atoms bonding with one carbon atom will give carbon a full outer shell of 8 and hydrogen full outer shells with 2 electrons each. This is now called methane (CH4.)

vi) Ammonia (NH3) 

Nitrogen has five outer electrons, and needs three more to become stable. It forms a bond with three hydrogen atoms, which each have one electron, to become stable. This is now known as ammonia (NH3).

vii) Oxygen (O2) 

In oxygen gas, one oxygen atom shares two pairs of electrons with another oxygen atom to form a double covalent bond. They each have six electrons in their outer shell and react with eachother, sharing two electrons each and becoming stable. 

viii) Nitrogen (N2)

Nitrogen atoms need three more electrons to become stable, so two nitrogen atoms can share three pair of electrons to fill their outer shells and create a triple bond. 

ix) Carbon dioxide 

Carbon atoms have four electrons in their outer shell and need another four to become full. Oxygen atoms each have six electrons in their outer shell and each share two of these with carbon, becoming stable. This has now formed two double covalent bonds and has become carbon dioxide (CO2).

x) Ethane (C2H6)

In ethane, six hydrogen atoms each share their only electron with one of the two carbon atoms. These two carbon atoms in return share their last electrons with eachother in a single covalent bond. This produces Ethane (C2H6).

xi) Ethene 

Here, four hydrogen atoms share their only electron with one of the two carbon atoms. In return, the two carbon atoms share their last two electrons with eachother to form Ethene, a carbon-carbon double bond. 

1.41 understand that substances with simple molecular structures are gases or

liquids, or solids with low melting points

The atoms within a molecule are held together by very strong covalent bonds, but the forces of attraction between these molecules are very weak. Because of the weak forces between the molecules, the melting and boiling points are very low as the molecules are easily parted from eachother. Therefore, most substances with simple molecular structures are gases, liquids or solids with low melting points. 

1.42 explain why substances with simple molecular structures have low melting

and boiling points in terms of the relatively weak forces between the

molecules

A substance with a simple molecular structure is one that only contains a few atoms in a molecule.The forces between these molecules are weak, and because of this it does not take much energy or heat to be supplied in order to break them up. They therefore have low melting and boiling points as not much heat or energy is needed to be supplied to break them up.

1.43 explain the high melting and boiling points of substances with giant covalent

structures in terms of the breaking of many strong covalent bonds

A substance with a giant covalent substance will have many atoms bonded together. The forces between these molecules will be very strong as there are so many of the bonds, and that means it will take a lot of energy or heat to break them up. Because of this, substances with giant covalent structures have very high melting and boiling points. 

1.44 draw diagrams representing the positions of the atoms in diamond

and graphite

Structure of diamond. Each carbon atom forms four covalent bonds in a very rigid giant covalent structure. It's because of this structure that diamond is the hardest natural substance in the world.

Structure of graphite. Each carbon atom forms three covalent bonds, creating layers which are free to slide over eachother. This leaves free electrons, which make for a good conductor of electricity, so graphite is the only non-metal which is a good conductor of electricity. 

1.45 explain how the uses of diamond and graphite depend on their

structures, limited to graphite as a lubricant and diamond in cutting.

Diamond is the hardest natural substance in the world as it has so many bonds held together in such a rigid giant covalent substance. Because of this, it is extremely hard and is therefore great for cutting as it could cut anything. 

In graphite, each carbon atom forms three covalent bonds which creates layers which are free to slide over eachother as the forces between them are so weak. This means that it is a slippery substance and can be used as a lubricant.