Inorganic Chemistry Flashcards
(169 cards)
1
Q
3.2.1
What is the definition of periodicity?
A
It is the quality or character of being periodic; the tendency to recur at intervals
2
Q
3.2.1
How is an element classified?
A
An element is classified as s,p,d,f block when the highest energy electrons are in an s,p,d,f sub-shell.
3
Q
3.2.1
What are the 2 physical properties trends of period 3 elements?
A
- Atomic radius
- melting points
4
Q
3.2.1
Why does atomic radius decrease across a period?
A
Atomic radius decreases across a period because…
* As the number of protons increase across the period, nuclear charge increase therefore stronger nuclear attraction.
* Therefore outer electrons (which are in the same level across the period) are drawn closer to the nucleus.
5
Q
3.2.1
Why does the melting point across a period increase?
A
The m.p.s of metallic elements increase across the period.
* Na ➜ Al, number of outer electrons increase
* Therefore more electrons can be delocalised, leading to a greater attraction between positive ions and delocalised electrons.
* Size of ions decrease across the period, leading to smaller ions, and therefore greater attraction between positive ions and delocalised electrons.
* Therefore more energy required to overcome attraction
6
Q
3.2.1
Why is the melting point of silicon very high?
A
The melting point of silicon is very high.
* Si has a giant covalent structure (compared to giant metallic for Na, Mg + Al).
* Therefore a lot of energy is required to break the strong covalent bonds.
7
Q
3.2.1
Why are the melting points of phosphorus, sulphur, chlorine and argon low?
A
The m.p.s of phosphorus (P4), sulphur (S8), chlorine (Cl2) + argon (Ar) are low.
* P4, S8 and Cl2 are simple covalent molecules and little energy is required to overcome the weak van der Waals’ forces between the molecules.
* Ar exists as atoms and very little energy is required to overcome the weak van der Waals forces between the atoms.
8
Q
3.2.1
Why do the melting points of elements after Si increase from phosphorus to sulfur then decrease again?
- attractions between molecules
- nuclear charge
A
The m.p.s of the elements after Si increase from phosphorus to sulfur then decrease again.
* Attractions between molecules = van der Waals’ forces (as molecules = non-polar).
* Greater the number of electrons, greater the induced dipole attractions, therefore greater van der Waals’ forces of attraction between molecules.
* This increases energy required to overcome attractions.
- P4 number of electrons < S8 no. of electrons therefore m.p. increase P4 to S8.
- S8 number of electrons > Cl2 no. of electrons and Ar has very low m.p. as monoatomic (resulting in very weak van der Waals forces) therefore m.p. decrease again.
9
Q
3.2.1
Why does the first ionisation energy increase across a period?
A
- Increase nuclear charge with similar shielding across the period, leading to stronger nuclear attraction.
- Therefore atomic radius decrease across the period.
- Therefore outer electron closer to the nucleus.
*More energy required to remove outer electron.
10
Q
3.2.1
Why are the first ionisation energy of atoms of group 3 elements lower than expected?
A
- Group 3’s outer electron is in the p sub shell whilst group 2’s outer electron is in the s sub-shell.
- Therefore group 3’s outer electron is further from the nucleus than group 2’s + has more shielding than group 2 (due to more inner electrons).
- Therefore weaker nuclear attraction, so electron more easily removed (less energy required to do so).
11
Q
3.2.1
Why are the first ionisation energy of atoms of group 6 elements is lower than expected?
A
- Group 6 atoms have a p4 arrangement the repulsion of 2 electrons in the same p orbital leads to less energy being required to remove the outer electron.
12
Q
3.2.2
What are the trends of atomic radius down Group 2?
A
Atomic radius increases down the group because…
*No. of shells increases down the group
*Therefore more shielding + therefore outer electrons further away from nucleus so…
*weaker nuclear attraction
13
Q
3.2.2
What are the trends of first ionisation energy down Group 2?
A
First ionisation energy decreases down the group because:
*Number of shells increases down the group
*Therefore more shielding and Therefore outer electron further away from nucleus so
*weaker nuclear attraction therefore outer electron held less tightly
*Therefore easier to remove outer electron (requires less energy).
14
Q
3.2.2
What are the trends of melting point down Group 2?
A
M.p.s decrease down the group because…
*Ions all have a +2 charge and same number of delocalised electrons per atom.
*However, size of metal ions increase down the group therefore atomic radius increase
*Therefore outer electrons further from nucleus so weaker nuclear attraction
*Therefore strength metallic bonding decrease
15
Q
3.2.2
What are the first 4 metals in Group 2?
A
- Magnesium (Mg)
- Calcium (Ca)
- Strontium (Sr)
- Barium (Ba)
16
Q
3.2.2
What occurs when Magnesium (Mg) reacts with water?
A
Reacts very slowly with warm water - few bubbles.
* Mg(s) + 2H2O(l) → Mg(OH)2 + H2 BUT, reacts vigorously with steam.
* Mg(s) + H2O(g) → MgO + H2
- MgO is formed rather than steam as the hydroxide is not stable at higher temperatures and thermally decomposes to give MgO and H2O.
17
Q
3.2.2
What occurs when Calcium (Ca) reacts with water?
A
Reacts with cold water and fizzing is seen in an exothermic reaction. White ppt. forms.
* Ca + 2H2O → Ca(OH)2 + H2
18
Q
3.2.2
What occurs when Strontium (Sr) reacts with water?
A
Reacts vigorously with cold water and fizzing is seen in a highly exothermic
reaction.
* Sr + 2H2O → Sr(OH)2 + H2
19
Q
3.2.2
What occurs when Barium (Ba) reacts with water?
A
Reacts violently with cold water and fizzing is seen in a highly exothermic reaction.
* Ba + 2H2O → Ba(OH)2 + H2
20
Q
3.2.2
What is the trend of reactivity down Group 2?
- Why is this the trend
A
Reactivity increases down the group because
* Metals lose outer electrons when they react.
* As we go down the group number of shells in therefore more shielding.
* Therefore weaker nuclear attraction and electron held less tightly by nucleus as we go down the group.
* Hence electron is lost more easily and metal = more reactive.
21
Q
3.2.2
How can the Relative Solubilities of Group 2 Hydroxides in Water be found?
- what must be added
A
The relative solubilities can be found by adding a solution of NaOH to solutions of the Group 2 ions + observing ppt.
22
Q
3.2.2
What is the solubility of Magnesium hydroxide?
A
Sparingly soluble in water (as sol. is slightly alkali indicating some OH- dissolved)
= thick white ppt.
* Mg2+(aq) + 2OH-(aq) → Mg(OH)2(s)
23
Q
3.2.2
What is the solubility of Calcium hydroxide?
A
Slightly soluble in water = white ppt.
* Ca2+(aq) + 2OH-(aq) → Ca(OH)2(s)
24
Q
3.2.2
What is the solubility of Strontium hydroxide / Barium hydroxide?
A
Soluble in water = no ppt.
25
3.2.2
What is the trend in solubility of hydroxides down Group 2?
The solubility of hydroxides down Group 2 increases
26
3.2.2
How can the Relative Solubilities of Group 2 sulphates in Water be found?
- what must be added
The relative solubilities can be found by adding a solution of Na2SO4 to solutions of the Group 2 ions + observing ppt.
27
3.2.2
What is the solubility of Mganesium sulfate?
Soluble in water = no ppt.
28
3.2.2
What is the solubility of Calcium sulfate?
Sparingly soluble in water = thin white ppt.
* Ca2+(aq) + SO42-(aq) → CaSO4(s)
29
3.2.2
What is the solubility of Stronstium sulfate?
Insoluble in water = white ppt.
* Sr2+(aq) + SO42-(aq) → SrSO4(s)
30
3.2.2
What is the solubility of Barium sulfate?
Insoluble in water = thick white ppt.
* Ba2+(aq) + SO42-(aq) → BaSO4(s)
31
3.2.2
What is the trend in solubility of sulphates down Group 2?
The solubility of sulphates down Group 2 decreases
32
3.2.2
How do you test for sulfate ions?
1) To 1cm3 of unknown solution add 1cm3 of dilute HCl acid.
* The HCl acid removes any other ions (e.g. carbonate ions) which may affect the test by giving a white ppt. with barium chloride solution = false positive result.
2) Add 1cm3 of barium chloride solution. If sulphate ions present: white ppt formed Ba2+(aq) + SO42-(aq) → BaSO4(s)
- If no sulphate ions present: no ppt formed
33
3.2.2
What are the Selected uses of the Magnesium and it's Compounds?
Magnesium is used in the extraction of titanium from its ore.
* The main titanium ore, titanium(IV) oxide (TiO2) is first converted to titanium(IV) chloride (TiCl4) by heating it with carbon in a stream of chlorine gas.
* The titanium chloride is then purified by fractional distillation, before being reduced by magnesium in a furnace at almost 1000°C
- TiCl4 + 2Mg → Ti + 2MgCl2
34
3.2.2
What are the Selected uses of the Magnesium hydroxide and it's Compounds?
Mg(OH)2 is used in medicine as a suspension in water known as ‘milk of magnesia’ to neutralise excess acid in the stomach.
35
3.2.2
What are the Selected uses of the Calcium hydroxide and it's Compounds?
Ca(OH)2 is used in agriculture in solid form known as ‘slaked lime’ to neutralise acidic soil.
36
3.2.2
What are the Selected uses of the Calcium oxide / Calcium carbonate and it's Compounds?
CaO / CaCO3 is used to remove SO2 from flue gases (= by-products of the combustion of fossil fuels) to prevent the formation of acid rain.
37
3.2.2
What are the Selected uses of the Barium sulfate and it's Compounds?
BaSO4 is used in medicine as a ‘Barium meal’ given to patients to eat who need x rays of their intestines.
* The barium absorbs the X-rays and so when BaSO4 gets to the gut the outline of the gut can be located using X-rays.
* Although Barium ions are toxic it is safe to use here because barium sulphate is insoluble therefore not absorbed into the blood.
38
3.2.3
What is the appearance of group 7 elements?
fluorine, chlorine, bromine and iodine
* Fluorine = poisonous pale yellow gas.
* Chlorine = poisonous green gas.
* Bromine = toxic red-brown volatile liquid - it forms a red-brown vapour.
* Iodine = shiny grey solid - it sublimes to form a violet vapour w/ gentle heating.
39
3.2.3
What are the trends in properties down group 7?
- Electronegativity decreases down the group
- B.p.s increase down the group
- Oxidising power decreases down the group
40
3.2.3
* Number of shells increase down the group therefore size of atoms increase so.
* bonding electrons in a covalent bond are further from the nucleus + there is more shielding.
* Therefore weaker attraction between positively charged nucleus and bonding pair of electrons.
* Therefore element has lower electronegativity.
41
3.2.3
* Size of diatomic molecules increase down the group.
* Larger molecules have more electrons, leading to greater induced dipole-dipole forces.
* Therefore greater van der Waals’s forces between molecules.
* Therefore more energy required to overcome the greater van der Waals’s forces as you go down the group.
42
3.2.3
An oxidising agent accepts electrons.
* Size of ions increase down the group.
* Therefore outer electrons are more shielded and further away from the nucleus.
* Therefore electrostatic force of attraction by nucleus on the additional electron becomes weaker down the group.
* Therefore harder to gain an electron.
43
3.2.3
in displacement reactions, what would we expect from our knowledge of oxidising powers of halogens decrease down the group?
…Fluorine to displace all other halogens from solutions of halide
compounds.
…Chlorine to displace bromide and iodide from solutions of bromide
iodide compounds.
…Bromine to displace iodide from a solution of an iodide compound.
44
3.2.3
What occurs when potassium chloride solution KCl (aq) - colourless - reacts with Cl2 (aq), Br (aq), I2 (aq)?
- Cl2 (aq) : no reaction
- Br (aq) : no reaction
- I2 (aq) : no reaction
45
3.2.3
What occurs when potassium bromide solution KBr (aq) - colourless - reacts with Cl2 (aq), Br (aq), I2 (aq)?
- Cl2 (aq) : orange solution (Br2) formed
- Br (aq) : no reaction
- I2 (aq) : no reaction
46
3.2.3
What occurs when potassium iodide solution KI (aq) - colourless - reacts with Cl2 (aq), Br (aq), I2 (aq)?
- Cl2 (aq) : brown solution (I2) formed
- Br (aq) : brown solution (I2) formed
- I2 (aq) : no reaction
47
3.2.3
What does a reducing agent do?
A reducing agent donates electrons. They, themselves, are OXIDISED.
48
3.2.3
What does a reducing power do?
- why does it increase down group 7
Reducing power increases down the group because:
* Size of the ions increase down the group.
* Therefore outer electrons are more shielded and further away from the
nucleus.
* Therefore electrostatic force of attraction by nucleus on outer
electrons becomes weaker down the group.
* Therefore easier to lose an electron.
49
3.2.3
What occurs in reactions of solid sodium halides with concentrated sulphuric acid?
- Fluoride (F-)
NaF + H2SO4 → NaHSO4 + HF NOT REDOX - it is an ACID/BASE REACTION
* Observations: misty white fumes (HF) are evolved.
50
3.2.3
What occurs in reactions of solid sodium halides with concentrated sulphuric acid?
- Chloride (Cl-)
NaCl + H2SO4 → NaHSO4 + HCl NOT REDOX - it is an ACID/BASE REACTION
* Observations: misty white fumes (HCl) are evolved.
51
3.2.3
What occurs in reactions of solid sodium halides with concentrated sulphuric acid?
- Bromide (Br-)
NaBr + H2SO4 → NaHSO4 + HBr NOT REDOX - it is an ACID/BASE REACTION
2HBr + H2SO4 → Br2 + SO2 + 2H2O REDOX
* Observations, reaction 1:
* misty white fumes (HBr) and red-brown vapour (Br2) are evolved.
* Observations, reaction 2:
- Br is oxidised from -1 (in HBr) to 0 (in Br2) therefore bromine = oxidation product. - S is reduced from +6 (in H2SO4) to +4 (in SO2) therefore sulphur dioxide = reduction product.
52
3.2.3
What occurs in reactions of solid sodium halides with concentrated sulphuric acid?
- Iodide (I-)
NaI + H2SO4 → NaHSO4 + HI NOT REDOX
2HI + H2SO4 → I2 + SO2 + 2H2O REDOX
6HI + H2SO4 → 3I2 + S + 4H2O REDOX
8HI + H2SO4 → 4I2 + H2S + 4H2O REDOX
* Observations (all these reactions happen in succession - not in isolation): misty white fumes evolved (HI), purple vapour evolved (I2), yellow solid formed (S), rotten egg smell (H2S) + black solid formed (I2).
* Redox:
➜ For the 1st redox reaction.
- the I is oxidised from -1 (in HI) to 0 (in I2) the S is reduced from +6 (in H2SO4) to +4 (in SO2)
➜ For the 2nd redox reaction,
- the I is oxidised from -1 (in HI) to 0 (in I2) the S is reduced from +6 (in H2SO4) to 0 (in S)
➜ For the 3rd redox reaction,
- the I is oxidised from -1 (in HI) to 0 (in I2) the S is reduced from +6 (in H2SO4) to -2 (in H2S)
Therefore oxidation product = iodine; reduction products = sulphur dioxide, sulphur and hydrogen sulphide.
53
3.2.3
How do you identify halide ions with dilute nitric acid (HNO3), followed by silver(I) nitrate (AgNO3) solution?
* Add dilute nitric acid (HNO3), followed by silver(I) nitrate (AgNO3) solution.
Ag+(aq) + Cl-(aq) → AgCl(s) = white ppt. formed
Ag+(aq) + Br-(aq) → AgBr(s) = cream ppt. formed
Ag+(aq) + I-(aq) → AgI(s) = yellow ppt. formed
- The dilute nitric acid removes other ions which would react with the silver nitrate solution (e.g. carbonates/sulphates/hydroxides)
N.B. silver nitrate solution does not form ppt. with fluoride ions in solution as silver fluoride is soluble in water. So, we can’t use it as test.
54
3.2.3
How do you identify halide ions with ammonia solution?
AgCl: will redissolve in dilute and concentrated ammonia solution to form a colourless solution.
AgBr: does not redissolve in dilute ammonia solution but does redissolve in conc. ammonia solution forming a colourless solution.
AgI does not redissolve in dilute or conc. ammonia solution.
55
3.2.3
What reaction occurs when chlorine mixes with water?
When you mix chlorine with water, it undergoes disproportionation (This means Chlorine, Cl2 is both oxidised and reduced).
Cl2 + H2O ⇌ HClO + HCl
* The Cl is oxidised from 0 (Cl2) to +1 (HOCl)
* The Cl is reduced from 0 (Cl2) to -1 (HCl)
56
3.2.3
What are the advantages + disadvantages of chlorine and chlorate?
- Chlorate(I) ions kill bacteria, this is why chlorine is used in water treatment to kill bacteria. It’s been used to treat drinking water and the water in swimming pools.
- The benefits to health (including, the irradiation of bacterial diseases like cholera) from using chlorine OUTWEIGH the fact it’s toxic / carcinogenic to humans.
In sunlight (UV), chlorine can decompose water to form chloride ions and oxygen.
2Cl2 + 2H2O ⇌ 4HCl + O2
57
3.2.3
What occurs in the reaction of Chlorine with Cold, Dilute, Aqueous NaOH (making bleach)?
- what is observed
When you mix chlorine gas with cold, dilute, sodium hydroxide solution at room temperature it undergoes disproportionation.
Cl2 + 2NaOH → NaCl + NaClO + H2O
* The Cl is oxidised from 0 (Cl2) to +1 (NaClO)
* The Cl is reduced from 0 (Cl2) to -1 (NaCl)
Observation: green gas forms a colourless solution.
One of the products formed is sodium chlorate(I), NaClO. It’s in solution in this case and that is bleach (which kills bacteria)
58
3.2.4
59
What are the general properties of transition metals?
Characteristics include:
* complex formation
* formation of coloured ions
* variable oxidation state
* catalytic activity
## Footnote
These properties arise from an incomplete d sub-level in atoms or ions.
60
Why is zinc not considered a transition metal?
Zinc can only form a +2 ion, which has a complete d orbital and does not have an incomplete d orbital in any of its compounds.
61
Define a complex in the context of transition metals.
A complex is a central metal ion surrounded by ligands.
62
What is a ligand?
A ligand is an atom, ion, or molecule that can donate a lone electron pair.
63
what is a lewis acid?
electron pair acceptor
64
what is a lewis base?
electron pair donor
65
ligands act as lewis_______? becasue?
bases because they donate a lone pair to form a coordinate bond
66
metal ions act as lewis______? because?
acids because they accept the lone pair
67
What is co-ordinate bonding?
Co-ordinate bonding is when the shared pair of electrons in the covalent bond comes from only one of the bonding atoms.
68
What is the co-ordination number?
The co-ordination number is the number of co-ordinate bonds formed to a central metal ion.
69
What are monodentate ligands? Give examples.
Monodentate ligands can form one coordinate bond per ligand. Examples include:
* H2O
* NH3
* Cl-
70
What are bidentate ligands? Give examples.
Bidentate ligands have two atoms with lone pairs and can form two coordinate bonds per ligand. Examples include:
* NH2CH2CH2NH2
* ethanedioate ion (C2O4^2-)
71
What are multidentate ligands? Give an example.
Multidentate ligands can form multiple coordinate bonds per ligand. An example is EDTA, which can form six coordinate bonds.
72
What happens during substitution reactions involving ligands?
Ligands like H2O, NH3, and Cl- can act as monodentate ligands, and exchange occurs without a change in co-ordination number.
73
What is the result of adding a high concentration of chloride ions to an aqueous transition metal ion?
It leads to a ligand substitution reaction and can involve a change in coordination number.
74
What occurs when solid copper chloride is dissolved in water?
It forms the aqueous [Cu(H2O)6]2+ complex, not the chloride [CuCl4]2- complex.
75
What is the coordination number change for Cu and Co when concentrated HCl is added?
The coordination number changes from 6 to 4.
76
Fill in the blank: A ligand substitution reaction can involve a change of _______.
[co-ordination number]
77
True or False: All transition metals can form coloured ions.
True.
78
What are bidentate ligands?
Ligands that can form two coordinate bonds to a metal ion.
79
Provide an example of a bidentate ligand.
Ethane-1,2-diamine
## Footnote
Ethane-1,2-diamine has the formula NH2CH2CH2NH2.
80
What is the coordination number of a complex with three bidentate ligands?
6
81
Write the equation for the formation of a complex with ethane-1,2-diamine.
[Cu(H2O)6]2+ + 3NH2CH2CH2NH2 → [Cu(NH2CH2CH2NH2)3]2+ + 6H2O
82
What is the shape and bond angle of a complex with bidentate ethanedioate ligands?
Octahedral shape with 90° bond angles.
83
What occurs during the partial substitution of ethanedioate ions in a copper(II) solution?
Four water molecules are replaced and a new complex is formed.
84
Write the equation for the formation of a complex with ethanedioate ligands.
[Cu(H2O)6]2+ + 2C2O4^2- → [Cu(C2O4)2]2- + 4H2O
85
What are multidentate ligands?
Ligands that can form multiple coordinate bonds to a metal ion.
86
Provide an example of a multidentate ligand.
EDTA4-
## Footnote
EDTA can form six coordinate bonds per ligand.
87
Write the equation for the formation of a complex with EDTA.
[Cu(H2O)6]2+ + EDTA4- → [Cu(EDTA)]2- + 6H2O
88
What is haem?
An iron(II) complex with a multidentate ligand.
89
what 2 does does a haem have?
Fe2+ centre
porphyrin ligand taking up four of the 6 coordination sites
90
How does oxygen interact with haemoglobin?
Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin.
91
True or False: Carbon monoxide (CO) can form a stronger coordinate bond with haemoglobin than oxygen.
True
92
Fill in the blank: The formula for the EDTA anion is _______.
C10H12N2O8^4-
93
What is the chelate effect?
The substitution of a monodentate ligand with a bidentate or multidentate ligand leads to a more stable complex.
94
How can the chelate effect be explained?
In terms of a positive entropy change as there are more molecules of products than reactants.
95
What is the reaction for the chelate effect involving copper and EDTA?
[Cu(H2O)6]2+ + EDTA4- → [Cu(EDTA)]2- + 6H2O
96
What happens to the number of moles in the copper and EDTA reaction?
Increases from 2 to 7, creating more disorder.
97
What is the enthalpy change in the chelate effect reaction?
Small, as there are similar numbers of bonds in both complexes.
98
What is the sign of free energy change (ΔG) when entropy change (ΔS) is positive and enthalpy change (ΔH) is small?
ΔG will be negative.
99
What are some applications of EDTA complexes?
* Removing poisonous heavy metal ions from rivers
* Included in shampoos to remove calcium ions from hard water
100
What is the reaction involving cobalt and ammonia in complex formation?
[Co(NH3)6]3+ + 3NH2CH2CH2NH2 → [Co(NH2CH2CH2NH2)3]2+ + 6NH3
101
What happens to the number of moles in the cobalt ammonia reaction?
Increases from 4 to 7.
102
What is the enthalpy change (ΔH) in the cobalt ammonia reaction?
Close to zero, as the number of dative covalent bonds remains the same.
103
In EDTA titrations, what is the ratio of EDTA to metal ions?
Always the same 1:1 ratio.
104
What is the initial concentration of the EDTA solution used in the titration?
0.0150 mol dm-3
105
How many cm³ of EDTA solution were required for the complete reaction in the river water sample?
6.45 cm³
106
How do you calculate the moles of EDTA used in the titration?
moles = conc x vol = 0.0150 × 6.45/1000
107
What is the concentration of copper(II) ions in the river water sample?
0.00387 mol dm-3
108
What shape do transition metal ions commonly form with small ligands?
Octahedral complexes
## Footnote
Examples of small ligands include H2O and NH3.
109
What shape can transition metal ions form with larger ligands?
Tetrahedral complexes
## Footnote
An example of a larger ligand is Cl-.
110
What is an example of a square planar complex?
Cisplatin
## Footnote
Square planar complexes are typically seen in certain transition metals.
111
What type of complexes does Ag+ commonly form?
Linear complexes
## Footnote
An example is [Ag(NH3)2]+, used as Tollen's reagent.
112
What are the two types of stereoisomerism shown by complexes?
Cis-trans isomerism and optical isomerism
113
Cis-trans isomerism is a special case of _______.
E-Z isomerism
114
What is the cis form of Ni(NH3)2Cl2 represented as?
Cis-Ni(NH3)2Cl2
115
What is the trans form of Ni(NH3)2Cl2 represented as?
Trans-Ni(NH3)2Cl2
116
What is an example of cis-trans isomerism in octahedral complexes?
Cis-[Cr(H2O)4Cl2]+
117
What is an example of optical isomerism in octahedral complexes?
Complexes with 3 bidentate ligands can form two optical isomers.
118
What characterizes optical isomers?
Non-superimposable mirror images
119
What changes can cause color changes in ions?
1. Oxidation state
2. Coordination number
3. Ligand
120
What color does [Co(H2O)6]2+ appear as?
Pink
121
What color does [CoCl]2+ appear as?
Blue
122
In the equation [Co(NH3)6]2+ → [Co(NH3)6]3+ + e-, what is changing?
Oxidation state
123
What color does the solution turn when [Co(NH3)6]2+ is formed?
Yellow
124
What color does the solution turn when [Co(NH3)6]3+ is formed?
Brown
125
How does color arise in transition metal complexes?
From electronic transitions from the ground state to excited states between different d orbitals
126
What happens to d electrons during color development?
A portion of visible light is absorbed to promote d electrons to higher energy levels
127
What is the average energy of d orbitals affected by?
The field of ligands
128
What does the equation AE = hv represent?
The energy difference between split d orbitals
129
What does 'v' represent in the equation AE = hv?
Frequency of light absorbed (unit s^-1 or Hz)
130
What is Planck's constant (H)?
6.63 × 10^-34 (J s)
131
What does a solution appear when it absorbs orange light?
Blue
132
What happens when a ligand or coordination number changes?
It alters the energy split between the d-orbitals, changing AE and the frequency of light absorbed
133
Why do Sc^3+ ions not exhibit color?
They have no d electrons left to move around
134
What is the electronic configuration of Zn^2+ that leads to no color?
3d10
135
What does spectrophotometry measure?
The amount of light absorbed by a colored complex ion
136
What is the relationship between light absorption and concentration?
The amount of light absorbed is proportional to the concentration of the absorbing species
137
What is the purpose of adding a suitable ligand in spectrophotometry?
To intensify the color of pale complexes
138
What is the method for determining the concentration of colored ions?
1. Add appropriate ligand
2. Make up solutions of known concentration
3. Measure absorption or transmission
4. Plot graph of absorption vs concentration
5. Measure absorption of unknown and compare
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What influences the redox potential for a transition ion?
The redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced by pH and by the ligand.
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What are the general trends in oxidation states of transition metals?
- Relative stability of +2 state with respect to +3 state increases
across the period
- Compounds with high oxidation states tend to be oxidising agents
e.g. MnO4
- Compounds with low oxidation states are often reducing agents e.g.
V2+ & Fe2+
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What are the successive oxidation states of Vanadium? solution colours
Vanadium has four main oxidation states
VO2+ Oxidation state +5 ( a yellow solution)
VO2+ Oxidation state +4 (a blue solution)
V3+ Oxidation state +3 (a green solution)
V2+ Oxidation state +2 (a violet solution)
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What happens if zinc is added to Vanadium in acidic solution?
Addition of zinc to the vanadium (V) in acidic solution will
reduce the vanadium down through each successive oxidation state, and the colour will successively change from yellow to blue to green to violet.
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why are redox titrations used?
metal changes oxidation state (gain or loss of electrons)
144
why is the solution acidic when metal ion reacts with water?
some hydrolysis of happens
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what makes the solution more acidic?
higher the charge on the metal in the metal aqua ion
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why are redox titrations used?
to measure the concentration of an oxidising or a reducing agent.
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what is the most common oxidising agent in redox titrations?
potassium manganate VII
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equation for the reduction of manganate ion during redox titrations
MnO4- + 8H+ + 5e- → Mn2+ + 4H2O
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KMnO4 needs to be in acidic solution for redox titrations, which acid is used?
DILUTE sulfuric
150
why cant you use HCl in redox titrations with KMnO4?
MnO4- would also oxidise Cl- to Cl2 so affect the volume of KMnO4 required in the titration
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why cant you use CONC sulfuric acid or CONC nitric acid in redox titrations with KMnO4?
they are oxidising agents themselves so affect the volume of KMnO4 required in the titration
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why cant ethanoic acid be used in redox titrations with KMnO4?
it is a weak acid and would not provide enough H+ ions.
153
what colour is potassium manganate?
purple
154
when potassium manganate reacts and forms Mn2+ what is the colour change?
purple to colourless
155
when is the end point of a redox titration using potassium manganate?
the first hint of pink and colour remains
156
what type of redox titration is it when potassium manganate is used?
self indicating
157
what is a self-indicating titration?
does not need an indicator added
158
half equation for the Fe in redox titration with potassium manganate
Fe2+ → Fe3+ + e-
159
overall equation of redox titration between Fe2+ and KMnO4
5Fe2+ + MnO4- + 8H+ → 5Fe3++ Mn2+ + 4H2O
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ratio of iron to manganate ions in redox titration
5:1
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ethanedioate ion
C2O4 2-
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half equation of ethanediote ion in redox titration with KMnO4
C2O42- → 2CO2 + 2 e-
163
overall equation of redox titration between ethanedioate ion and KMnO4-
5C2O42- + 2MnO4- + 16H+ → 10CO2 + 2Mn2+ + 8H2O
164
why is the reaction between MnO4- and C2O42- slow to begin with? how would you do this reaction as a titration?
ions are both negative so they repel each other
165
as the C2O42- and MnO4- ions repel and reaction is slow to start with, what must be done in this titration?
needs warming
166
what is a catalyst?
Catalysts increase reaction rates without getting used up. They do this by providing an alternative route with a lower activation
167
what can transition metals and their compounds act as? two types of catalysts?
Transition metals and their compounds can
act as heterogeneous and homogeneous
catalysts.
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