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. 

Ionic Compounds (Section 1f)

Ionic Compounds 

1.28 describe the formation of ions by the gain or loss of electrons
In order for atoms to become stable, they must have a full outer shell. To do this, they join with other atoms and either lose or gain electrons to have a full outer shell. When an atom has lost or gained electrons, it has become a charged particle, or an ion. 

1.29 understand oxidation as the loss of electrons and reduction as the gain of

electrons

Oxidation is the process of an atom losing electrons to become a positive ion (cation.) 
Reduction is the process of an atom gaining electrons to become a negative ion (anion.)

OIL RIG = Oxidation is loss, Reduction is gain. 


1.30 recall the charges of common ions in this specification

Hydrogen = H+ 
Lithium = Li+ 
Sodium = Na+ 
Potassium = K+ 

Beryllium = Be 2+
Magnesium = Mg 2+
Calcium = Ca 2+ 

Boron = B 3+ 
Aluminium = B 3+ 

Fluorine = F- 
Chlorine = Cl - 
Bromine = Br - 

1.31 deduce the charge of an ion from the electronic configuration of the atom

from which the ion is formed

The electronic configuration of the atom will tell you how many electrons it has in it's outer shell. From this, you know whether to round up or down. For example, sodium has an electronic configuration of; 2.8.1. This means that it has one electron in it's outer shell and will therefore need to get rid of one electron in order to become stable. Therefore, the electronic configuration of the sodium ion will be 2.8 and the charge will be; Na +1, as it has lost one electron and is therefore a positive ion (cation.) 

1.32 explain, using dot and cross diagrams, the formation of ionic compounds by
electron transfer, limited to combinations of elements from Groups 1, 2, 3

and 5, 6, 7
Dot and cross diagrams represent the process of two atoms bonding to form ions. One atom will have it's electrons displayed as dots, and the other as crosses so it is clear what electrons are being gained/lost in each atom.

This diagram shows what happens when sodium (Na) reacts with Chlorine (Ch) to form sodium chloride. Sodium has one electron in it's outer shell, and gives this to chlorine which has seven electrons in it's outer shell, so they both become stable. The red arrow indicates the electron that the sodium atom loses, and the chlorine atom gains. 

1.33 understand ionic bonding as a strong electrostatic attraction between

oppositely charged ions

In ionic bonding, atoms lose or gain electrons to have a full outer shell and become stable. Once an atom has lost or gained electrons, it has become an ion.One of the atoms will become a positive ion, and the other a negative. These ions will then be strongly attracted due to the electrostatic attraction between ions with opposite charges ( + and -).

1.34 understand that ionic compounds have high melting and boiling points

because of strong electrostatic forces between oppositely charged ions.

In order to melt or boil anything, heat must be applied to break bonds. The stronger the bond is, the more heat will need to be applied. Because ionic compounds have such strong bonds, a significant amount of heat is needed for them to melt or boil, hence why they have very high melting and boiling points. 

1.35 understand the relationship between ionic charge and the melting and boiling point of an ionic compound

The charges of an ion affect the strength of the ionic bonding. The higher the charge, the stronger the bond. For example, a lattice of 3+ and 3- ions will have a stronger force of attraction between them than a lattice of 1+ and 1- ions. This means that the bigger the difference in charge (the larger the numbers,) the higher the melting and boiling point of the ionic compound. 

1.36 describe an ionic crystal as a giant three-dimensional lattice
structure held together by the attraction between oppositely

charged ions

When an ionic compound is formed, the positively charged ions (cations) attract the negatively charged ions (anions) and arrange themselves into a three-dimensional lattice structure, held together by the attraction between the ions.

1.37 draw a diagram to represent the positions of the ions in a crystal of

sodium chloride.

The diagram shows the arrangement of the ions in a crystal of sodium chloride.