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A2 chemistry 4.4
Edexcel chemistry - entropy
| Question | Answer |
|---|---|
| What is the first law of thermodynamics? | Energy can't be made or destroyed |
| What is entropy? What factors increase it? | Entropy is a measure of the disorder of a system and of the number of ways of distributing the energy of a substance in the available energy levels, J/K/mol. Factors increasing: gas<liquid<solid, high temperature, more gas moles, more random arranegemnt |
| What does the standard entropy of a substance depend on? | Physical state and complexity (number of atoms) |
| What are the 3 equations for entropy + precuations? | ΔStotal = ΔSsystem + ΔSsurroundings. ΔSsystem = ΣSproducts - ΣSreactants (including stoichiometric numbers). ΔSsurroundings = -ΔHsystem/T. T must be in K - celcius + 273. ΔH kJ/mol -> J/mol |
| What is the 2nd law of thermodynamics? | Spontaneous change results in an increase in disorder/entropy. Natural direction of change is of increasing total entropy. 1 ordered arrangement (crystalline substance) an be rearranged into many disordered arrangements, so P(disorder) is greater |
| What is a spontaneous reaction? | Occurs without continuous intervention, although can be briefly heated/ catalysed |
| State ΔH, ΔSsystem, ΔSsurroundings, effect of increasing/ decreasing T for exothermic reaction | -ΔH, - ΔSsystem, + ΔSsurroundings, increasing T makes ΔH/T less positive so ΔStotal is less positive |
| State ΔH, ΔSsystem, ΔSsurroundings, effect of increasing/ decreasing T for endothermic | +ΔH, +ΔSsystem, -ΔSsurroundings, increasing T makes ΔH/T less negative so ΔStotal is more positive |
| What is the 3rd law of thermodynamics? | Entropy of a perfect crystalline substance at 0K is 0, but entropy is gained as soon as heat/ defect is introduced into system |
| Explain temperature/ entropy graph | 0 entropy @ 0K. At melting+boiling points, large increase in S as change in state allows particles to move more randomly +freely. No T change -energy is used to break bonds. Supplying heat to hot object smaller increase in S than same to cold object. |
| Explain how kinetics and thermodynamics affect reaction feasibility | Kinetics - determines rate, large Ea means slow reaction because reactants are more kinetically stable than products, thermodynamically feasible reaction may not take place at RT due to high Ea. THermodyanmics- enthalpy = exothermic is favourable... |
| Cont. | Entropy: ΔStot>200 reaction goes to completion, 0<ΔStot<200 equilibrium favours RHS, -200<ΔStot<0 equilibrium favours LHS, ΔStot<-200 reaction doesn’t go. ΔStot>0 reaction is thermodynamically feasible + spontaneous, products more stable than reactants |
| When can endothermic reactions occur at RT? Give some examples + equations | If the magnitude of ΔSsys > ΔH/T. NH4NO3 (s) -> NH4+ (aq) + NO3- (aq). 2CH3COOH + (NH4)2CO3 -> CH3COONH4 + CO2 + H2O. Ba(OH)2.8H2O (s) + NH4Cl (s) -> 10H2O + BaCl2 + 2NH3 |
| Why is the combustion of magnesium ribbon in air considered spontaneous? | Exothermic reaction, the increase in entropy of the surroundings compensates for the decrease in ΔSsys (gas -> solid). Since ΔS total > 0, thermodynamically feasible and spontaneous (2nd law thermodynamics). Continues to burn when heat is removed |
| Define ΔH soln + implications for ionic solids | The enthalpy change when 1 mol of solid dissolves in sufficient solvent to give an infinitely dilute solution, ions become hydrated when ionic solids dissolve in water. |
| Define ΔH hyd | The enthalpy change when 1 mol gaseous ions dissolves in sufficient solvent to give an infinitely dilute solution, usually exothermic |
| What factors make ΔH hyd more exothermic? | High charge density on ions (higher charge, smaller ionic radius) means stronger attraction between ions and solvent molecules, more bonds form releasing energy |
| What is ΔH LE? | Enthalpy change when 1 mol of ionic solid is formed from gaseous constituent ions that start infinitely far apart, exothermic |
| What factors make ΔH LE more exothermic? | Higher charge density (higher charge on ions, smaller ionic radii) = stronger electrostatic attraction between ions. More polarisation of anion - polarising cation (high charge density) + polarisable anion (high charge, large radius) = more covalency |
| What equation allows you to calculate theoretical values of ΔH soln? Give an example for CaCl2 | ΔH soln = Δ H hyd - ΔH LE. e.g: ΔHsoln (CaCl2) = ΔHhyd (Ca2+) + 2ΔHhyd (Cl-) - ΔHLE (CaCl2) |
| When is ΔH soln exothermic? | ΔHhyd is exothermic enough to compensate for large positive (-ΔHLE) i.e. energy released from making bonds between solvent+solute compensates for energy to break lattice. Ions have a high charge density: form strong attractions with solvent, lots of energ |
| When is ΔH soln endothermic? | If magnitude of ΔHLE> ΔHhyd, so the energy released from forming bonds between solvent/solute can't compensate for energy required to break lattice. Tend to be ions with LOW CHARGE DENSITY- form weaker attractions, less energy released by hydration |
| Give an example of an ionic solid that dissolves endothermically and explain why it occurs at RT | NaCl: low charge density, weak attractions between ions and solvent, less energy released by hydration so endothermic, but increase in Ssys compensates for this: ionic solid -> ions + solvent doesn't become much more ordered |
| What determines the extent of solubility of a solid? | ΔS total for 1 mol of solid |
| Does solubility increase or decrease down a group? | ΔH hydration and ΔH LE become less exothermic down a group- if ΔHhyd becomes less exothermic more quickly, solubility decreases down group. If ΔH LE becomes less exothermic more quickly, solubility increases down group |
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