Chemistry Unit 5 Definitions & Common Questions

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Enthalpy of formation
The enthalpy change when 1 mole of a compound is formed from its elements under standard conditions, all substances being in their standard states
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Ionisation enthalpy
The minimum amount of energy required to remove one mole of electrons from 1 mole of atoms in the gaseous state
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Enthalpy of atomisation
The enthalpy change for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
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Bond dissociation enthalpy
The enthalpy change to break the bond in 1 mole of gaseous molecules to form gaseous atoms
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Electron affinity
The enthalpy change when one mole of gaseous atoms forms one mole of gaseous negative ions
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Second electron affinity
The enthalpy change when one mole of electrons is added to one mole of gaseous negative ions to form ions each with a doubly negative charge
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Lattice dissociation enthalpy
The enthalpy change to separate 1 mole of an ionic substance into its gaseous ions
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Lattice formation enthalpy
The enthalpy change when one mole of a solid ionic compound is formed from its gaseous ions
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Enthalpy of hydration
The enthalpy change when one mole of gaseous ions is converted into aqueous ions
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Enthalpy of solution
The enthalpy change when one mole of solid dissolves in water to form aqueous ions
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Enthalpy of combustion
The enthalpy change when one mole of a compound is completely burned in oxygen under standard conditions, all substances being in their standard states
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Mean bond enthalpy
Enthalpy change for breaking a covalent bond in the gaseous state, averaged over a range of compounds
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O symbol in ΔHfO
Standard conditions
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Perfect ionic model
Ions can be regarded as perfect spheres
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Amphoteric
Act as acids and bases
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Oxidation
Loss of electrons
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Reduction
Gain of electrons
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Oxidising agent
Gains electrons
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Reducing agent
Loses electrons
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Lewis acid
Electron pair acceptor
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Lewis base
Electron pair donor
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ΔE
Energy difference between d-orbitals
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h
Planck's constant
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v
Frequency of light absorbed
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Heterogeneous
In a different state from reactants
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Homogeneous
In the same state as the reactants
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Autocatalyst
The catalyst is a reaction product
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Ligand
Electron pair donor
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Unidentate
Donates one electron pair
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Bidentate
Donates two electron pairs from two different atoms
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Multidentate
Forms two or more co-ordinate bonds
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Complex
A transition metal ion co-ordinately bonded to one or more ligands
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Co-ordinate bond
Where one atom provides both electrons for the formation of a covalent bond
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Co-ordination number
The number of co-ordinate bonds formed to a central metal ion
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Chelate effect
The substitution of a unidentate ligand by a multidentate ligand giving a more stable complex and an increase in entropy
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Predict whether the lattice enthalpy of dissociation that you have calculated from the Born-Haber Cycle will be less than, equal to or greater than the value obtained from a perfect ionic model.
Greater. Covalent character makes the ionic bond stronger so it requires more energy to break
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Explain why the lattice dissociation enthalpy of magnesium chloride is greater than that of calcium chloride
The magnesium ions are smaller than calcium ions. The attraction between magnesium ions and chloride ions is stronger and requires more energy to break.
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Explain why the lattice dissociation enthalpy of magnesium oxide is greater than that of magnesium chloride.
The oxide ion has a greater charge. The attraction between the magnesium and oxide ions is stronger and requires more energy to break.
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Suggest why the 2nd electron affinity is an endothermic process
The negative ion repels the added electron
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Suggest why the electron affinity of chlorine is an exothermic change
Due to the attraction between the nucleus of chlorine and the added electron
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Explain why this reaction is feasible at all temperatures
ΔG= ΔH - TΔS, ΔH is negative and ΔS is positive, hence ΔG is always negative
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Explain, in terms of structure and bonding, why sodium oxide has a high melting point
Na2O has a giant ionic lattice structure, with strong electrostatic forces of attraction between oppositely charged ions,which require lots of energy to break
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Explain why phosphorus oxide has a higher melting point than sulphur (V1) oxide
P4O10 is a bigger molecule than SO3, so there are increased and stronger Van der Waals forces between molecules
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Explain why sulphur dioxide has a low melting point
It has a simple molecular structure, with weak Van der Waals forces between molecules which don't require much energy to break
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Explain why silicon dioxide had a high melting point
It has a macromolecular structure, with stong covalent bonds between atoms which require lots of energy to break
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Write an equation for the reaction that occurs between Na2O and P4O10 in the absence of water. State the general type of reaction illustrated by this example
6Na2O + P4O10 --> 4Na3PO4, acid-base reaction
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Write an equation for the reaction between aluminium oxide and an excess of aqueous sodium hydroxide
Al2O3 + 2NaOH + 3H2O --> 2NaAl(OH)4
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Write an equation for the reaction between silicon dioxide and sodium hydroxide
2NaOH + SiO2 --> Na2SiO3 + H2O
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Identify a solution that could be used in the salt bridge. Give 2 reasons why this solution is suitable for its purpose.
KNO3, doesn't interfere with the redox reaction, allows the movement of ions
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Explain why the ammeter reading would fall to 0 after a time
The reactants are used up
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Give one reason why the cell cannot be electrically recharged
The reaction is not reversible
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Give one reason why the e.m.f. of the lead-acid battery changes after several hours
The reactants are used up
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Explain why the voltage remains constant in a fuel cell
There is a continuous supply of fuel, the concentrations remain constant
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Suggest 2 properties of platinum that make it suitable for use as an external circuit in the cell
It conducts electricity, and is unreactive/inert
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Describe a standard hydrogen electrode
Give components and conditions: Hydrogen gas, 1.0 moldm-3 HCl, at 298K and 100kPa, Pt electrode
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What are the standard conditions which apply to Fe3+(aq)/Fe2+(aq) when measuring this potential?
Both Fe3+(aq) and Fe2+(aq) have a concentration of 1.0 moldm-3
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Suggest, in terms of electrons, why the colours of the complex ions are different
The energy levels of the d electrons are different, there is a different energy difference (ΔE) between d orbitals, so a different frequency of light is absorbed when d electrons are excited and a different frequency of light is transmitted.
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In terms of electrons, explain why aqueous iron (II) ions are green
The d orbitals split in energy levels, electrons are promoted to higher energy d orbitals by absorbing energy in the form of light, the colour observed is a combination of the colours not absorbed
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Give the equation that relates the energy change ΔE to Planck's constant, h and the frequency of the visible light,v
ΔE= hv
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State 3 different features of transition metal complexes that cause a change in the value of ΔE, the energy change between the ground state and the excited state of the d electrons
Oxidation state, type of ligands, number of ligands, shape of complex ion
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Explain, in terms of the chelate effect why the [Co(EDTA)]2- is formed in preference to [Co(H2O)6]2+
The number of moles increases (from 2 to 7), therefore disorder increases, so the entropy change is positive, and a more stable complex is formed
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State the general property of transition metals that allows the vanadium in vanadium oxide to act as a catalyst in the Contact Process.
Variable oxidation states
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Write 2 equations to show how vanadium oxide acts as a catalyst in the Contact Process.
V2O5 + SO2 --> V2O4 + SO3, V2O4 + 1/2O2 --> V2O5
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Write the overall equation for the Contact Process
SO2 + 1/2O2 --> SO3
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The oxidation of C2O42- ions by MnO4- ions in acidic solution is an example of a reaction that is autocatalysed. Write 3 equations to show how the autocatalyst is formed and how it is involved in this oxidation of C2O42- ions.
1) 2MnO4- + 5C2O42- + 16H+ --> 2Mn2+ + 10CO2 + 8H2O 2) MnO4- + 8H+ + 4Mn2+ --> 5Mn3+ + 4H2O 3) 2Mn3+ + C2042- --> 2Mn2+ + 2CO2 (1 + 3 can be derived from 1/2 eqns - remember C2O42- --> CO2 and MnO4- --> Mn2+, 2 needs learning)
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Write out the equations and observations for the Fe2+, Co2+, Cu2+. Fe3+, Cr 3+. Al 3+ species as: 1) aqua ions 2) with NaOH 3) with excess NaOH 4) with dilute NH3 5) with excess/conc. NH3 6) with Na2CO3 7) with conc. HCl
Check your notes to see if correct. Helpful tips: 1) check if ion is a 3+/2+ ion - 2+ ions form carbonates, 3+ don't, 2) compounds beginning with c undergo NH3 subtitution (partially if Cu) 3) neutral= ppt charged= soln
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Give the observations when the Period 3 elements (Na - S) react w O2 to form their oxides
Na= bright yellow flame, Mg= bright white flame, white ash, Al= burns w bright light, P4= white smoke, limited O2 P2O3 forms, red+ white allotropes, S= blue flame, choking gas
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Write the equations and observations for the reduction of Cr2O72- to Cr(H2O)63+ and then to Cr(H2O)62+ by Zn.
1) 3Zn + 14H+ + Cr2O72- --> 2Cr3+ + 3Zn2+ +7H2O Observation: orange solution to green solution 2) 2Cr3+ + Zn --> 2Cr2+ + Zn2+ Observation: green solution to blue solution
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Write the equation and observations for the acid base equilibrium for chromium in its oxidation state of 6+
2CrO42- + 2H+ ⇌Cr2O72- + H20 Observations: CrO42-= yellow chromate ions, Cr2O72= orange dichromate ions
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Write the overall equation and separate equations for the reduction of S2O82- ions to SO42- ions by I-
Overall: S2O82- + 2I- --> 2SO42- + I2, 1) S2O82- + 2Fe2+ --> 2SO42- + 2Fe3+ 2) 2Fe3+ + 2I- -> 2Fe2+ + I2 (Remember: I- is oxidised overall, needs to react with Fe3+ which can be reduced and oxidise I-, steps can occur in either order)
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Give the formula and draw the structure of a linear complex
[Ag(NH3)2]+
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Give the formula and draw the structure of a tetrahedral complex
[CoCl4]2- (Cl- ions too big to fit more than 4 around)
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Give the formula and draw the structure of a square planar complex
[Pt(NH3)Cl2] (neutral complex, usually Ni and Pt form square planar complexes)
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Give the formula and draw the structure of an octahedral complex
[Fe(H2O)6]2+/[Cr(NH3)6]2+ (there are loads more examples, usually octahedral if ligand=H2O/NH3)
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Draw out the structure of cis-platin and trans-platin
Remember cis is the same side, both Cls above, Trans= across, one up one down Check: http://www.ch.ic.ac.uk/local/projects/s_liu/Html/Graphics/Cisplatin.gif , http://patentimages.storage.googleapis.com/WO2000051600A1/imgf000003_0001.png
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Write out the equations for oxidising chromium [Cr(H2O)6]3+
Step 1: Add NaOH [Cr(H2O)6]3+ + 6OH- -> [Cr(OH)6]3- + 6H2O Step 2: Add H2O2 (oxidising agent) 2[Cr(OH)6]3- + 3H2O2 -> 2CrO42- + 2OH- + 8H20
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Write out the equation for the oxidation of cobalt in alkali solution
2[Co(OH)6]4- + H2O2 -> 2[Co(OH)6]3- + 2OH-
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Write out the equations for the oxidation of cobalt in ammonical solution
Step 1 (form neutral ppt): [Co(H2O)6]2+ + 2OH- -> [Co(OH)2(H2O)4] + 2H2O Step 2: dissolve in excess NH3 [Co(OH)2(H2O)4] + 6NH3 -> [Co(NH3)6]2+ + 4H2O + 2OH- Step 3: (oxidise using H2O2 or O2) H2O2 + 2[Co(NH3)6]2+ -> 2OH- + 2[Co(NH3)6]3+
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Write out half equations for the reduction of manganate ions, the oxidation of iron(II) ions, the reduction of dichromate ions, and the reduction of H2O2
1) MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O 2) Fe2+ -> Fe3+ + e- 3) Cr2O72- + 14H+ + 6e- -> 2Cr3+ + 7H2O 4) H2O2 + 2e- -> 2OH-
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Give the equation and a catalyst for making methanol
CO (g) + 2H2 (g) -> CH3OH (g) (Heterogeneous Cu(s)/ZnO(s)/Al2O3(s) catalyst)
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Give the equation and catalyst for the Haber Process
1/2 N2 (g) + 3/2 H2 (g) ⇌NH3 (g) (Heterogeneous Fe(s) catalyst)
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Write out the half equations and overall equation for the hydrogen fuel cell
1) H2 -> 2H+ + 2e- 2) 4H+ + 4e- + O2 -> 2H2O 3) Overall: 2H2 + O2 -> 2H2O
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Give some of the disadvantages of the hydrogen fuel cell
H2 is flammable, water is a greenhouse gas, source of H2=crude oil, electrolysis of H2O requires electricty (burns fossil fuels, releases greenhouse gases)
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Card 2

Front

Ionisation enthalpy

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The minimum amount of energy required to remove one mole of electrons from 1 mole of atoms in the gaseous state

Card 3

Front

Enthalpy of atomisation

Back

Preview of the front of card 3

Card 4

Front

Bond dissociation enthalpy

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Preview of the front of card 4

Card 5

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Electron affinity

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