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Chemistry II Exam Review

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CHEM 1020U
Omer Yukseker

Chemistry II Exam Review Part One: Thermodynamics 1. Thermodynamics so far… - Conservation of Energy: Energy cannot be created or destroyed. It can only be converted from one form to another. - State Function: function or property whose value depends on the present state or condition of the system, not the pat used to arrive at that state. 2. First Law of Thermodynamics - Energy is conserved, not created 3. Enthalpy (H) - H = E + PV - At constant pressure, the enthalpy change is equal to the heat, ΔH = q p 4. Second Law of Thermodynamics - The entropy of the universe always increases 5. Entropy (s) - The measure of the amount of disorder in a system - ΔS = ΔS + ΔS > 0 universe system surroundings - This determines which reactions will occur - If Δsystem< 0 for a reaction, the value of Δsurroundingsst be positive, and be larger than ΔS system - The reaction will NOT occur if this is the case - ΔS = ΣS products – ΣS reactants 6. Third Law of Thermodynamics - A perfect crystal at zero kelvin is assigned an entropy of zero. - Suggests that there is two sources of disorder in a system: - i. A perfect crystal - ii. At zero Kelvin - Perfect Crystal: there is disorder due to the positions of atoms, molecules, etc. - Ssolid Sliquid<<< S gas - At zero kelvin: there is disorder due to the movement of atoms, molecules etc. - In a given phase, the entropy will go up with T. - Kelvin scale: it starts at absolute zero below which there is no colder temperature than 0K 7. Gibb’s Free Energy - G ≡ H – TS - ΔG = ΔH – Δ(TS) - For Isothermal reactions we use: ΔG = ΔH – TΔS (at constant temperature) - At constant T, P, the sign of ΔG tells us whether a reaction is spontaneous or not. - Until now, the only way to determine spontaneity was to calculate Δs universe - Calculating ΔG is much easier, since its value depends only on the system. - As the temperature increases, the reaction will go from spontaneous to not so spontaneous - ΔG° rxn= ΣΔG°(products) - ΣΔG°(reactants) ΔH ΔS ΔG Spontaneity (Constant T and P) <0 >0 Always <0 Always spontaneous <0 <0 <0 at low temperature Spontaneous at low temperatures >0 >0 <0 at high temperature Spontaneous at high temperatures >0 <0 Always >0 Never spontaneous Part Three: Equilibrium 1. Chemical Equilibrium - We saw in kinetics that the rate of a reaction depended on the height of the energy barrier (E )a - A larger Eameans a slower reaction, since a lower percentage of all molecules will have enough energy to overcome the barrier - What about the reverse reaction? - products reactants - The activation energy for the reverse reaction (Ea,ris rate  rate larger than the activation energy for the forward forward reverse reaction a b c d k fA] [B]  k [C] [D] r - the reverse reaction should occur more slowly - Until now we have only considered reactions where c d E is so large that the effective rate of the reverse k f [C] [D] a,r  a b  K reaction is zero k r [A] [B] - rate = k [A] 1 forward forward 1 - rate reverse= kreverseB] - When the forward and reverse reactions happen at the same rate, we have reached CHEMICAL EQUILIBRIUM - The concentrations of chemicals do not change - their rates of production and loss are equal - since the chemistry (the reactions) are ongoing (they don’t stop), this is called DYNAMIC EQUILIBRIUM - aA + bB cC + dD - the forward rate = kf[A] [B] b c d - the reverse rate = k rC] [D] 2. Aspects of Equilibrium States - They display no macroscopic evidence of change - They are reached through spontaneous processes - They show a dynamic of forward and reverse processes - They are the same regardless of direction of approach 3. Reaction Quotient - Q = [C] [D] /[A] [B] a b - If Q = K, then we know the system is at equilibrium - What happens if Q ≠ K? then the system is not at equilibrium, Q must change until Q = K - If Q < K, then we know that the system is not at equilibrium - System shifts towards products, [C] and [D] must increase and [A] and [B] must decrease - System shifts towards reactants, [A] and [B] must increase and [C] and [D] must decrease 4. Equilibrium & Thermodynamics - ΔG = ΔG° + RT ln Q - At equilibrium, Q = K, and ΔG = 0. So, 0 = ΔG° + RT ln K – ΔG° = RT ln K ln K = – ΔG°/RT – ΔG°/RT K = e 5. Temperature Dependence of K - at constant T, ΔG° = ΔH° – T ΔS° - so , K = e – (ΔH°/RT – TΔS°/RT) – (ΔH°/RT – ΔS°/R) K = e ln K = ΔH°/RT –ΔS°/R 6. Upsetting Equilibrium - Starting with a system at equilibrium, say one (or more) of the following happens: - reactants are added - products are added - solution is diluted - gasses are compressed/expanded - in these cases, Q will change, and Q ≠ K - temperature is changed - in this case, K will change, and Q ≠ K - In all of these cases, the concentrations/pressures will change to re-establish equilibrium 7. Le Chatalier’s Principle - A system brought out of equilibrium will shift to re-establish equilibrium - Say A B is at equilibrium, what happens if more of B is added? - System will shift to the forward reaction to move back to equilibrium Part Four: Organic Chemistry 1. What is Organic Chemistry? - Study of carbon compounds - Carbons bond together in short chains, long chains, and rings - Methane to DNA - Covalent Bonds (Polar Covalent Bonds) - Mulitple Bonds (Double or Triple Bonds) - VSEPR models (for shapes of molecules) - “Organic chemistry” now refers (roughly) to the chemistry of C,H,O,N and P compounds 2. Hybridization of Carbon 3 - sp (tetrahedral) 1 s and 3 p atomic orbitals - sp (trigonal planar) 1 s and 2 p atomic orbitals - sp (linear) 1 s and 1 p atomic orbitals 3. Structures and Properties - The structures of molecules have a huge influence on their behaviour, properties - we can look at a struc
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