Notesale: Turn your study into money

Document Preview

Extracts from the notes are below, to see the PDF you'll receive please use the links above

---

Unit 5 Chemistry

Useful formulas

Enthalpy change of
solution
Enthalpy change of
solution

Bond enthalpies
Measuring enthalpy
change
Enthalpy change
Entropy change
Gibbs free energy
Energy difference orbitals/
subshells

Thermodynamics Page 1

12
...
Lattice formation enthalpy:

2
...
2 Born-Haber cycles
20 December 2015

19:26

Thermochemical cycle
• Includes all the enthalpy changes
• Formation of an ionic compound
• All elements in standard states have 0 enthalpy
Left hand side = right hand side
Group 1 elements

4
...


Atomisation
enthalpy of
chlorine

2
...


6
...


Lattice formation
enthalpy of NaCl

Enthalpy of
formation of
NaCl

Group 2 elements

5
...


first ea of clorine

First IE of Magnesium

3
...


Atomisation enthalpy of
magnesium

2
...


Enthalpy of formation of
MgCl

Chlorine
• 2 chlorines are involved
• 2 first electrin affinity
Magnesium
Thermodynamics Page 3

7
...
3 More enthalpy changes
21 December 2015

19:29

Enthalpy of solution
• Ionic solids dissolve well in polar solvents
• To dissolve- lattice must be broken up - requires input of energy
(lattice enthalpy)
• Separate ions are solvated by solvent molecules, water
• Clusters round the ions so the positive ions and the negative are
surrounded by the positive ends of dipoles
• Hydration - water used as a solvent

Dissolving an ionic compound in water
1
...
Endothermic - energy has been out in
2
...
Experimental value is bigger than calculated value
2
...
Ionic bonds- quite strongly polarised- covalent character

Thermodynamics Page 6

12
...
18 g-1k-1

q= heat change

+ kJmol-1 Endothermic
─ kJmol-1 Exothermic

Thermodynamics Page 7

12
...

Sodium reacts vigorously with cold water
• Fizzes forming a molten ball on the surface and produces H2 gas
pH= 13-14
• Magnesium reacts very slowly with cold water

Weak alkali is formed
pH=9-10

• Magnesium reacts faster with steam

Most period 3 elements readily react with oxygen
4Na(s) + O2(g) → 2Na2O(s)

Vigorous

Si(s) + O2(g) → SiO2(s)

Slow

4P(s) + 5O2(g) → P4O10(s)

Spontaneously
combusts

Na2O - bright
yellow flame
2Mg(s) + O2(g) → MgO(s)

Vigorous
MgO-

Limited supply of O2
White fumes Allotropes
White light
2P(s) + 1
...


Non-metal oxides
• Silicon oxide (macromolecular)
• Giant structure , strong covalent bonds and
high mp

Distortion electron cloud

• Phosphorus and sulphur oxides (molecular)
• IMF is van der Waals forces and dipole- dipole
forces
• Low mp
Na2O

MgO

Al2O3

SiO2

P4O10

SO3

SO2

Tm/K

1548

3125

2345

1883

573

290

200

Bonding/
structure

Giant
Ionic

Giant
Ionic

Giant
Ionic/
covalent

Giant covalent Simple
(macromolecul Covalent
ar)
(molecular)

Simple
Simple
Covalent
Covalent
(molecular) (molecular)

Basic oxides

• Produces strong alkaline solution

• Produces a somewhat alkaline solution

pH=14
pH=10
Acid oxides
Phosphorus pentoxide

• Reacts violently with water
• Produces acidic solution of Phosphoric acid

P4O10(s) + 6H2O(l) → 4H3PO4(aq)
pH= 3

Sulphur dioxide

SO2(g) + H2O(l) → H2SO3(aq)

• Reacts with water, soluble
• Produces sulphurous acid
• Partially dissociates

pH=2-3

Sulphur trioxide

SO3(g) + H2O(l) → H2SO4(aq)

• Reacts violently with water to produce sulphuric(VI) aid

Element

Na

Mg

pH=0-1

Al

Si

P

S

Formulae of
oxides

Na2O

MgO

Al2O3

SiO2

P4O10

SO2
SO3

Acid-base
character of

Basic
Strong

Basic
Somewhat

Amphoteric

Acidic

Strongly
Acidic

Acidic

Period 3 Page 10

character of
oxide

Strong
alkaline

Somewhat
alkaline

pH when
dissolved in
water

13 - 14

10

7

7

0-1

Weak 2 - 3
(SO2)
Strong 0-1
(SO3)

Solubility

Soluble

Sparingly
soluble

Insoluble

Insoluble

Soluble

Soluble

Acidic

Period 3 Page 11

Reactions with acids & bases
14 November 2015

19:33

Acid + base → salt + water

Basic oxides neutralise acids

Amphoteric reacts with acids and bases

Na2O (s) + 2HCl (aq) → 2NaCl (aq) + H2O (l)

Al2O3 (s) + 3H2SO4 (aq) → Al2 (SO4) (aq) + 3H2O (l)

MgO (s) + H2SO4 (aq) → MgSO4 (aq) + H2O (l)

Al2O3 (s) + 6HNO3 (aq) → 2Al(NO3)3 (aq) + 3H2O (l)
Al2O3 (s) +2NaOH + 3H2O (l) → 2NaAl(OH)4 (aq)

Sodium hydrogensulphate(IV)

Acid oxides neutralises bases

• Sulphur dioxide (weak acid) will react with a strong base
• Hot conc NaOH
• Colourless solution of sodium silicate
SiO2 (s) + 2NaOH (aq) → Na2SiO3 + H2O

Phosphorus pentoxide

•
•
•
•

Sulphur dioxide
SO2 (aq) + NaOH (aq)→NaHSO3 (aq)
NaOH (aq) + NaHSO3 (aq) → Na2SO3 (aq) + H2O (l)
Sodium sulphate(IV)

Phosphoric acid has 3 OH groups
Triprotic acid
React with sodium hydroxide in 3 stages
Phosphoric acid loses an H atom

H3PO4 (aq) + NaOH (aq) → NaH2PO4 (aq) + H2O (l)

NaH2PO4 (aq) + NaOH (aq) →Na2HPO4 (aq) + H2O (l)
Na2HPO4 (aq) + NaOH (aq) →Na3PO4 (aq) +H2O (l)

Overall reaction : H3PO4 (aq) + 3NaOH (aq) →Na3PO4 (aq) + 3H2O (l)

Period 3 Page 12

Redox equations
20 November 2015

16:38

OIL RIG

Oxidation Is Loss of electrons Increase in oxidation number
Reduction Is Gain of electrons Decrease in oxidation number

Oxidation agent Accepts electrons
Reducing agent Donates electrons

Molecules
Sum of the oxidation numbers = 0

Ion
The sum of the oxidation numbers = charge on the ion
Method 1: (half equations)

1
...
insert the number of electrons being gained or lost:
3
...
balance Oxygen atoms by adding H2O

5
...

Potential difference
• connect two different electrodes
• Cannot be measured directly
Salt bridge
• avoids any further metal/ion potentials
• Filter paper soaked in a solution ( usually with KNO3)

Negative electrode

Positive electrode

Oxidation

Reduction

Oxidised

Zn → Zn2+ + 2e-

Reduced

Cu2+ + 2e- → Cu

ee-

Overall: Zn + Cu2+ → Cu + Zn2+

Solution of its own ions

Standard 1 mol dm -3 Hydrogen electrode

Standard electrode potential of a half cell
• Voltage measured
• Standard conditions
• Half-cell is connected to the standard
hydrogen electrode

Electrode potential = 0
Always shown on the left

Salt bridge
Oxidation
Redox Page 14

Reduction

Salt bridge
Oxidation

Reduction

Phase boundary ( between solid and solution)

Redox Page 15

Predicting redox reactions
20 November 2015

17:22

1
...
More reactive a non-metal the more it wants to gain electrons
to form negative ion
3
...
Oxidising agent - The more positive half-cell will reduce ( gain
electrons) - higher oxidation number

Easily reduced

Easily oxidised

Reaction is feasible
• Actually happens
• Positive value

Redox Page 16

Electrochemical cells
20 November 2015

17:50

Non rechargeable cells
Zinc/ copper cells
Oxidised

Zn → Zn2+ + 2e-

Reduced

Cu2+ + 2e- → Cu

Overall: Zn + Cu2+ → Cu + Zn2+

• Electrons are transferred from the
more reactive metal to the least
reactive metal
• Voltage = potential difference

Porous pot (Semi permeable membrane) - allows ions to flow
completing the mixture of solutions
...
5-4V

NiO(OH) + H2O + e- → Ni(OH)2 + OHemf = 1
...
2V

General properties
28 November 2015

18:00

Transition metals

Electronic configurations

4s orbital usually fills up first

Chromium
• 1 electron each 3d sub shell
• 1 electron in 4s subshell
• More stable

Copper
• Full 3d sub shell
• One electron in the 4s sub shell
• More stable

Transition metal
• can form one or more stable ions
• Partly filled d-subshell
Physical properties
• High density
• High melting and boiling points
• Similar atomic radii

Transition metal Ions
• Loose their 4s electrons first
Chemical properties
•
•
•
•

Variable oxidation states
They form coloured ions
Good catalysts
Form complex ions

Transition metals Page 20

Complex ions and shape
28 November 2015

18:36

Complex ion
• One or more ligand attached to a central metal ion
• Dative covalent bond/coordinate bond

Ligand
• Donates a lone pair of electron to a transition metal
• Dative covalent bond/coordinate bond
Co-ordination number
• Number of co-ordinate bonds to the Ligand
• Surrounded by a d-block metal ion
Shape

Diagram

Tetrahedral

Co-ordination Ligands
number
4

Larger Ligands
• Smaller
complexes
• Cl-

octahedral

6

Small ligands
• H2O
• NH3

Linear

2

Ag+

Square
planner

4

CN-

multidentate ligands
• Have more than one lone pair
Transition metals Page 21

• Have more than one lone pair
• Bond to a transition metal
EDTA4+ (Ethylenediaminetetraacetic
acid)
• Acts as a hexa-dentate ligands
• Lone pairs 4 oxygen atoms and 2
nitrogen atoms
• Complex ion with polydentate
ligands- chelates
• Effective in removing a d-block
metal from solution

Bi-dentate ligands

Ethane-1,2-diamine
• Lone pair electrons on both
nitrogen atoms
• Name is shortened to en
• [Cr(en)3]3+
• Neutral ligand

Benzene- 1,2-diol
• Neutral ligand

Ethanedioate (oxalate )
ion C2O32-

Transition metals Page 22

Coloured ions
29 November 2015

10:51

Partly filled d subshell

is energy difference of orbitals/subshells (J)
h= planks constant(JS)

Ligands split the 3d subshell in to TWO energy levels

v= frequency of light absorbed (Hz = s-1)

• Certain wavelengths/ frequency of light absorbed
• d electron can move from one orbital to next
• d electron exiled
• Absorb energy in visual spectrum = difference
between energy levels
• Absorbs light from the visual spectrum
• We see the colours which are not absorbed

Colorimetry - find the ratio of metal ions to ligands in a complex

Colour of metal complexes
• Depends on the oxidation state of the metal
• Also on the ligands (and therefore the shape of the
complex ion)
• Different compounds of the same metal will have
different colours

1
...

3
...


Transition metals Page 23

Mix the two solutions ( metal ion and ligand in different proportions)
White light is shone through the filter and is only colour of light is let through
Light passes through the sample to the light detector( detects absorbance)
Concentration increases , more light absorbed, less light will pass through the
solutions

Variable oxidation states
29 November 2015

11:18

Transition metals Page 24

15
...
3 colours complex ions and ppt
05 December 2015

09:26

little OH- or
NH3

Ion

Excess OH-

Excess NH3

Conc HCL

[Fe(H2O)6]2+ Fe(H2O)4(OH)2 Fe(H2O)4(OH Fe(H2O)4(OH)2
)2
Pale green
Green ppt
Green ppt
Green ppt
solution
[Co(H2O)6]2+ Co(H2O)4(OH)2 Co(H2O)4(OH [Co(NH3)6]2+
)2
Pink solution Blue ppt
Blue ppt
Straw
coloured
[Cu(H O) ]2+ Co(H O) (OH) Co(H O) (OH solution) (H O
[Cu(NH
2

6

Plale blue
solution

2

4

Pale blue ppt

2

2

4

)2
Pale blue
ppt

3 4

2

CO32FeCO3
Green ppt

[CoCl4]2-

CoCO3

Blue
solution

Pink ppt

[CuCl4]2-

CuCO3

Yellow
solution

Blue- green ppt

]2+

)2
Deep blue
solution

[Fe(H2O)6]3+ Fe(H2O)3(OH)3 Fe(H2O)3(OH Fe(H2O)3(OH)3
)3
Yellow
Brown ppt
Brown ppt Brown ppt
solution ]3+ Al(H O) (OH) [Al(H O) (O Al(H O) (OH)
[Al(H2O)6
2 3
3
2 2
2 3
3
H)4]Colourless
White ppt
Colourless
White ppt
solution
solution

Brown ppt :Fe(H2O)3(OH)3

[Cr(H2O)6]3+

Cr(H2O)3(OH)3 [Cr(OH)6]3-

[Cr(NH3)6]3+

Green ppt :Cr(H2O)3(OH)3

Violet
solution

Green ppt

Purple
solution

+ CO2 bubles

Green
solution

lewis acids and bases Page 26

+ CO2 bubles

white ppt :Al(H2O)3(OH)3
+ CO2 bubles

16

---