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    Example: When ice melts DH = + 6.0 kJ mol-1. In terms of the intermolecular forces present comment on the sign of DH.

    Remember that bond breaking is an endothermic process so the process described above is endothermic since hydrogen bonds between the molecules need to be broken for the ice to melt.
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    http://filestore.aqa.org.uk/subjects...5-QP-JUN12.PDF

    can someone please explain 2(b) to 2(d) for me please

    i dont understand how Gibbs free energy change is represented on a graph
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    Periodicity - Study of the reaction of period 3 elements Na - Ar to illustrate periodic trends - Topic 2.

    Okay so here are the equations for the reactions of Na and Mg with water.

    2Na(s) + 2H2O(l) --------------> 2NaOH(aq) + H2(g)

    Mg(s) + 2H2O(l) ----------------> Mg(OH)2(aq) + H2(g)


    Na is a violent reaction with cold water whereas Mg reacts only very slowly with cold water.

    However Mg does react violently with steam to produce the white solid MgO:

    Mg(s) + H2O(l) ---------------> MgO(s) + H2(g)
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    Reaction of the elements of period 3 (Na to S) with oxygen.

    Okay, so I'll just write down the equations and state the observations for each element that we need to know (there are 6):

    ~

    Element: Sodium.

    Equation with Oxygen: 4Na(s) + O2(g) -----------> 2Na2O(s)

    Observations: Vigorous reaction, yellow flame, white fumes condense to white powder.

    ~


    Element: Magnesium.

    Equation with Oxygen: 2Mg(s) + O2(g) --------> 2MgO(s)

    Observations: Vigorous reaction, bright white light, white fumes condense to white powder.


    ~


    Element: Aluminium.

    Equation with Oxygen: 4Al(s) + 3O2(g) ----------> 2Al2O3(s)

    Observations: Harder to ignite, but bright white light, white powder


    ~

    Element: Silicon.

    Equation with Oxygen: Si(s) + O2(g) -------------> SiO2(s)

    Observations: Hard to ignite, but bright white light.


    ~

    Element: Phosphorus.

    Equation with Oxygen: 4P(s) + 5O2(g) ---------------> P4O10(s)

    Observations: Burns easily, white light, copious white fumes which condense to give white powder

    ~

    Element: Sulphur.

    Equation with Oxygen: S(s) + O2(g) ----------> SO2(g)

    Observations: Burns easily and slowly with a blue flame, pungent (choking) gas.

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    Okay, just a small extra part of info: The trend in these reactions is in the formula of the oxides with the oxidation state of the oxides increasing by 1. (except for SO2)
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    Periodicity - A survey of the acid base properties of the oxides of period 3 elements - Topic 2.

    Okay so the structure + bonding of each of the highest oxides of the elements Na – S in relation to their melting points:


    Na2O, MgO and Al2O3 all have ionic bonding with the ions held in a giant lattice structure.

    To melt these oxides the strong electrostatic forces of attraction between the oppositely charged ions holding the lattice together have to be broken requiring a lot of energyand giving rise to high melting points.


    ~


    SiO2 has covalent bonding in a macromolecular structure.

    To melt this oxide many strong covalent bonds have to be broken which requires a lot of energy giving rise to a high melting point.


    ~


    P4O10(s), SO2(g)and SO3(l)all have covalent bonding and molecular structures.·

    To melt these oxides weak intermolecular dipole-dipole or van der Waals’ forces between the molecules have to be brokenwhich requires little energy giving rise to low melting points.
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    The reactions of the oxides of the elements Na – S with water.

    Alright so, I'm going to write out the equations for the reactions of the oxides with water + state the PH of the resulting solutions.


    Compound: Sodium Oxide.

    Equation of oxide with water: Na2O(s) + H2O(l) ------------> 2NaOH(aq) (Strong base formed)

    PH: 14

    ~


    Compound: Magnesium Oxide.

    Equation of oxide with water: MgO(s) + H2O(l) -----> Mg(OH)2(aq) (reaction occurs to a limited extent forming a weak base)

    PH: 9.


    ~


    Compound: Aluminium Oxide.

    Equation of oxide with water: NO REACTION BC INSOLUBLE IN WATER.

    PH: 7


    ~

    Compound: Silicon dioxide.

    Equation of oxide with water: NO REACTION BC INSOLUBLE IN WATER.

    PH: 7.


    ~

    Compound: Phosphoric Acid.

    Equation of oxide with water: P4O10(s) + 6H2O(l) -------------> 4H3PO4(aq) (Strong triprotic acid formed)

    PH: 0.

    ~

    Compound: Sulphur trioxide.

    Equation of oxide with water: SO3(l) + H2O(l) ------------> H2SO4(aq) (Strong diprotic acid formed)

    PH: 0.

    ~


    Compound: Sulphur dioxide.

    Equation of oxide with water: SO2(g) + H2O(l) ----------> H2SO3(aq) (weak acid formed)

    PH: 3.
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    The trend in the reactions of the oxides with water in terms of the type of bonding present.


    Okay, so as the bonding changes from ionic to covalent, the solutions of the oxides change from alkaline to acidic.


    Basically, in general, metal oxides (ionic) are basic. (react with acids)

    AND

    non-metal oxides (covalent) are acidic. (react with bases).


    However Al2O3(s) is amphoteric and this means that it is capable of reacting with both acids and bases.
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    Keep going anon!
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    Okay, so now I'm just going to write out full equations as well as ionic equations for the reactions of the oxides with the acids or bases indicated.

    REMEMBER THE RULES THAT WE'VE JUST WRITTEN ABOVE:

    IONIC = REACT WITH ACIDS.

    COVALENT = REACT WITH BASES.

    ~

    Sodium oxide and sulphuric acid:

    Na2O + H2SO4 ----------> Na2SO4 + H2O

    Na2O(s) + 2H+(aq) ----------> 2Na+ (aq) + H2O (l)



    Magnesium oxide and hydrochloric acid:

    MgO + 2HCl ----------> MgCl2 + H2O

    MgO(s) + 2H+(aq) --------> Mg2+(aq) + H2O(l)



    Aluminium oxide and nitric acid:

    Al2O3 + 6HNO3 ---------> 2Al(NO3)3 + 3H2O

    Al2O3(s) + 6H+(aq) ---------> 2Al3+(aq) + 3H2O(l)



    Aluminium oxide and sodium hydroxide:

    Al2O3 + 2NaOH + 3H2O ----------> 2NaAl(OH)4

    Al2O3(s) + 2OH-(aq)+ 3H2O(l) → 2Al(OH)4 (Aluminate ion)


    (As you can see, Al reacts with both acid + bases like we said above)


    Silicon oxide and potassium hydroxide:

    SiO2 + 2KOH ----------> K2SiO3 + H2O

    SiO2(s) + 2OH(aq) → SiO32-(Silicate ion) + H2O(l)



    Phosphorus(V) oxide and sodium hydroxide:

    P4O10 + 12NaOH ---------> 4Na3PO4 + 6H2O

    P4O10(s) + 12OH(aq) ----------> 4PO43-(aq)(Phosphate ion) + 6H2O(l)



    Sulphur trioxide and potassium hydroxide:

    SO3 + 2KOH -----------> K2SO4 + H2O

    SO3 + 2OH(aq) --------> SO42-(aq)(Sulphate ion) + H2O(l)
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    I forgot to include some important information in the post before last so I'll just put it here:


    - Ionic oxides are basic because the O2-ions accept H+ ions to form OH- ions.

    - Covalent oxides are acidic because the non-metal (Si, P, S) is electron deficient and is attacked by :OH- ions.
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    Redox Equilibria - Redox equations - Topic 3.

    Oxidation - electron loss·

    Oxidising agent - species which gain/accept electrons.

    Okay so I'm just going to do some examples of half equations in acidic conditions + alkaline conditions.

    Firstly acidic:


    i) S2O82- ----------> SO42-

    therefore S2O82- + 2e ----> 2SO42



    ii) VO2+ -------------> V3+

    therefore 4H+ + VO2+ + 2e -----> V3+ + 2H2O



    iii) MnO4 ------------> MnO2

    therefore 4H+ + MnO4 + 3e ------> MnO2 + 2H2O



    iv) Cr2O72- --------> Cr2+

    therefore 14H+ + Cr2O72-+ 8e -----> 2Cr2+ + 7H2O
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    Now alkaline conditions:

    i) O2 --------> OH

    therefore 2H2O + O2 + 4e ----------> 4OH



    ii) H2 --------> H2O

    therefore 2OH + H2 -----------> 2H2O + 2e



    iii) [Cr(OH)6]3- -------> CrO42-

    therefore 2OH + [Cr(OH)6]3- ------------> CrO42- + 4H2O + 3e



    iv) Fe(OH)2 -------------> Fe(OH)3

    therefore OH+ Fe(OH)2 -----------> Fe(OH)3 + e
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    Cell diagram for when Mg2+(aq)/Mg(s)half-cell is coupled with the Fe3+(aq)/Fe2+(aq) half-cell with an e.m.f of 3.50 V and taking into account that the Mg electrode is negative:


    Mg(s) | Mg2+(aq)|| Fe3+(aq), Fe2+(aq) | Pt(s) E = +3.50 V


    Things to remember:

    - The most reduced forms are ALWAYS on the outside, (Mg + Fe2+ are the most reduced in this scenario)

    - '|' represents the separates phases (different states)

    - Pt electrode needed for Fe3+(aq)/Fe2+(aq) half-cell.

    - the + sign of electrode potential gives sign of the RH electrode.

    - The cell can also be written the other way round:


    Pt(s) | Fe2+(aq), Fe3+(aq)|| Mg2+(aq) | Mg(s) E = -3.50 V
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    Chemical substances that can be used in a salt bridge and the function of a salt bridge:


    The salts KNO3 or KCl are often used. The ions need to be free to move so are in saturated solution or a gel.

    A salt bridge maintains electrical neutrality in the cell by providing ions to each half-cell.
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    Redox Equilibria - Electrode potentials - Topic 3.

    The standard hydrogen electrode.

    Alright, so I'm just going to draw it below then I'll describe the diagram afterwards.
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    Posted from TSR Mobile
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    lol at that drawing. :'3

    Anyway, explanation of the diagram above:


    - Solution of hydrochloric acid of concentration 1.00 mol dm-3·

    - Hydrogen gas bubbled into solution at a pressure of 100kPa

    - Platinum electrode

    - Temperature = 298 K

    - Electrode reaction: H+(aq) + e- = ½H2(g)

    - Half-cell defined as 0V

    - Cell diagram: Pt(s) | H2(g) | H+(aq)
 
 
 
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