1 / 49

Energetics

Energetics. 5.1 Endothermic and Exothermic Reactions. Endothermic rxn  heat is taken in from the surroundings ( rxn vessel gets cooler) Exothermic rxn  rxn that result in the release of heat ( rxn vessel gets warmer)

megan-ware
Download Presentation

Energetics

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Energetics

  2. 5.1 Endothermic and Exothermic Reactions • Endothermicrxn heat is taken in from the surroundings (rxn vessel gets cooler) • Exothermic rxn rxn that result in the release of heat (rxn vessel gets warmer) • Enthalpy change amount of heat energy taken in/ given out in a chemical rxn (ΔH) • Positive for endo. Rxn • Negative for exo. rxn

  3. Exothermic Rxn

  4. Endothermic Rxn

  5. Stability • Exo. produces a more stable product • Ex. Cdiamond CgraphiteΔH= -1.9 kJ mol-1 • Kinetic vs. Thermodynamic stability • Graphite is more thermodynamically more stable than diamond • Diamond is kinetically stable • What does this mean?

  6. Activation Energy Define Activation Energy

  7. Activation energy • High AE= faster or slower rxn? • Can the conditions be altered? • If yes, in what ways? • Does how endo- or exothermic a rxn is tell us how quickly the rxn will run? • Ex. Which rxn will run faster? • -52 kJ mol-1 • -16 kJ mol-1 • Do endo. or exo. rxns usually occur spontaneously? (under normal conditions)

  8. 5.2 Calculations of enthalpy changes from experimental data • Specific heat capacity (c) energy required to raise the temp. of 1 g of substance by 1 K (1°C) or, the energy to raise 1 kg of substance by 1 K. • Units-> J g-1 K-1 J g-1 °C-1 kJ kg-1 K-1 J kg-1 K-1

  9. Calculating c • q= mcΔT • q= heat energy • m= mass • ΔT= change in temperature • How does heat capacity affect how easily a substance can be heated? • Can c be calculated for a substance undergoing cooling?

  10. Measuring enthalpy change of combustion rxns • Worked ex. Page 185 • Literature value for ΔH ethanol is -1371 kJ mol-1 • What are some reasons the calculated value is different? • A bomb calorimeter could have been used so that the system was heavily insulated and provide a plentiful supply of oxygen

  11. Enthalpy changes in solution • General method for measuring • Measure known amounts of reagents • Record initial temps • Mix in a polystyrene cup • Record max/ min temperatures observed • Assume that c for the final solution is the same as water

  12. ΔH of Solutions:Definitions • Enthalpy change of neutralisation (ΔHn) enthalpy change when 1 mol of water molecules are formed when acid reacts with alkali under standard conditions • H+(aq) + OH-(aq)  H2O(l) • Enthalpy change of solution (ΔHsol) the enthalpy change when 1 mol of solute is dissolved in excess solvent to form a solution of ‘infinite dilution’ under standard conditions • NH4NO3(s) NH4+(aq) + NO3-(aq)

  13. 5.3 Hess’s Law • The enthalpy change accompanying a chemical reaction is independent of the pathway between the initial and final states • What does this mean? • Ex. Find ΔHr for the reaction of AB • Knowns: • ΔHr= ΔH1 +ΔH2 • AC =ΔH1 • BC= ΔH2 • What is C? • What needs to change about the BC step?

  14. Hess’s law:Definitions • State function pathway does not matter • Standard conditions pressure= 1 atm (or, 1.01E5 Pa), 298K (or, 25°C) • Standard enthalpy change (ΔHrΘ) the enthalpy change when molar amounts of reactants as shown in the stoichiometric equation react together under standard conditions to give products (Θ= under standard conditions)

  15. Working out enthalpy changes • Hess’s law can be used to determine enthalpy changes of unknows from knowns • P. 194 Worked example

  16. 5.4 Bond enthalpies • The enthalpy change when 1 mole of covalent bonds, in a gaseous molecule, are broken under standard conditions (aka bond energy) • Ex. Then enthalpy of H-H bond is 436 kJ mol-1 H2(g)  2H(g) ΔHΘ= +436 kJ mol-1 How many H-H bonds were broken?

  17. Bond Enthalpy • What state must a substance be in to calculate bond enthalpy? • Consider this process: Br2(l)  2Br(g) Br-Br= 193 kJ mol-1 What is the ΔHΘ? +224 kJ mol-1 Why is this higher than the bond enthalpy? The reactants are not in a gaseous state We must also account for the energy required for vaporisation of the reactants This process is called atomisation

  18. Bond breaking • Bond breaking is… endothermic or exothermic? Endothermic! What does that mean about ΔH? Positive! What will bond making be? Exothermic with a negative ΔH

  19. Average bond enthalpy • The average amount of energy required to break 1 mole of covalent bonds, in a gaseous molecule under standard conditions • These are the values used to calculate bond enthalpies

  20. Using bond enthalpies to work out enthalpy changes in a rxn • Must draw out the structural formulas for rxn • Imagine the rxn happening and ALL bonds being broken • Add up the total energy for all broken bonds • Draw in all the bonds formed in products • Add up the total energy of all bonds made • Determine signs for the total enthalpy changes • Broken positive • Made negative • Add the changes to get the overall enthalpy change of the rxn

  21. Example • Consider the rxnbetween ethene and bromine, to produce 1,2-dibromoethane, C2H4(g) + Br2(g) C2H4Br2(g) What bonds are broken? What bonds are made? Follow your steps!

  22. Using a cycle • Same concept as previous calculations, but a process is drawn out to see all the steps • P. 203

  23. HL2

  24. 5.5 Calculating enthalpy changes:Definitions • Standard enthalpy change of combustion (ΔHcΘ) the enthalpy change when 1 mole of a substance is completely burnt in oxygen under standard conditions. • If ΔHcΘis always negative, what does this mean? • Standard enthalpy change of formation (ΔHfΘ) the enthalpy change when 1 mole of the substance is formed from its elements in their standard states under standard conditions • Endo and exorxns are dependent on the type of substance • ΔHfΘ for any element in its standard state is zero

  25. Using ΔHcΘto calculate enthalpy change • Method 1: Construct an enthalpy cycle • P. 208 • Method 2: rearrange the equations to give the overall equations related to the enthalpy change • P. 209 • Method 3: use an enthalpy level diagram for calculations • P. 210 • Method 4: use the equation, ΔHr = ΣΔHc (reactants)- ΣΔHc(products)

  26. Using ΔHfΘto calculate other enthalpy chages • Method 5: similar to method 1, but used for formation rather than combustion • P. 214 • Method 6: refer method 2 (be sure equations are running in the correct direction) • Method 7: draw enthalpy level diagram for formation (method 3) • Method 8: use the equation, ΔHr = ΣΔHf (products)- ΣΔHf (reactants)

  27. Choosing your method • Choose a method based on the data you are given, NOT on what needs to be found • If needing the enthalpy of combustion and given the enthalpy of formation, use one of the methods 5-8 • Once the basic principle of the methods are understood, there is no need to have any distinctions between them

  28. 5.6 Enthalpy changes for ionic compounds • First ionisation energy • Second ionisation energy • First electron affinity enthalpy change when one electron is added to each atom in 1 mol of gaseous atoms under standard conditions (always EXOTHERMIC) X (g) +e-  X- (g) • Second electron affinity (always ENDOTHERMIC) why? • Lattice enthalpy(ΔHΘlatt) the enthalpy change when 1 mol of an ionic compound is broken apart into iest constituent gaseous ions under standard conditions

  29. Born-Haber cycles • Enthalpy level diagram breaking down the formation of an ionic compound into a series of simpler steps • 1. put the equation for the enthalpy of formation • 2. add lattice enthalpy • 3. convert to gaseous form (why?) • Two steps • Must convert ALL reactants to gaseous form • Connect the cycle by adding the electrons removed from one reactant to the more electronegative reactant

  30. Draw a Born-Haber cycle • P. 219-221 • Na and Cl example

  31. Comparisons of lattice enthalpy • P. 224 • What is lattice enthalpy the result of? Electrostatic attractions of + and – ions • If the attractions of great, will more or less energy need to be supplied to break the bonds? More

  32. Effect of charge and size • How does the charge of the ions effect lattice enthalpy? The higher the ion charge, the greater the lattice enthalpy • Does NaCl or MgCl2 have great lattice enthalpy? MgCl2 • How does size effect lattice enthalpy? The larger the ions the weaker the forces, the smaller the lattice enthalpy • Which has the larger lattice enthalpy, CsCl or NaCl? NaCl

  33. Theoretical vs. experimental • Theoretical assumes a totally ionic model • What is this? Bonding is solely due to attractive forces between oppositely charged ions • Experimental use the Born-Haber cycle to find • These are compared to determine how ionic a particular compound is

  34. How to use theoretical and experimental values • If values are exactly the same, complete ionic bonding is suggested • If values are significantly different, it is suggested that the bonding has a significant degree of covalency • Ex. Silver iodide Theoretical value/ kJ mol-1 736 Experimental value/ kJ mol-1 876 What do the values suggest?

  35. Covalent character • What is covalent character the result of? Polarisation of the negative ion by the positive one • How does size of the anion effect this? The polarisation effect is greater

  36. 5.7 ENtropY • A measure of randomness or disorder of a system • Especially significant in the case of endothermic processes occurring at standard conditions (ice melting at room temp, water evaporating, NaCl dissolving in water etc) • Endo rxns can only occur if there is an increase in entropy • Represented by S • Units: J K-1 mol-1

  37. Standard entropy • Represented by SΘ • Positive ΔSΘ indicates increased entropy • Less order • Negative ΔSΘ indicates decreased entropy • More order

  38. Predicting sign of entropy change • Which state of matter has the higher entropy? Gases • Which state has the least entropy? Solids

  39. Period 2 entropies

  40. Predicting sign • Must consider whether the system’s disorder increases or decreases • Good to consider whether moles of gas have increased or decreased • What would an increase in moles of gas mean? • If moles of gas remain constant, our prediction of a change in entropy would be approximately zero

  41. Calculating entropy change for a rxn ΔSΘ= ΣSΘproducts -- ΣSΘreactants

  42. 5.8 Spontaneity • Spontaneous reaction one that occurs without any outside influence

  43. Predicting spontaneity • A reaction being spontaneous does not mean it will run quickly! • Whether a rxn is spontaneous or not under a certain set of conditions can be deduced by looking at the change in the “entropy of the Universe” ΔSUniverse = ΔSsuroundings + Δssystem If ΔSUniverseis positive, the entropy of the universe increases and the rxn occurs spontaneously When heat is given out in a rxn, the entropy of the surroundings get hotter

  44. Gibbs free energy • Represented by ΔG • Also called free energy change • ΔG = ΔH – TΔS • ΔH and ΔS are referring to the system • Under standard conditions this symbol is used: ΔGΘ • For a reaction to be spontaneous, ΔG for the rxn must be NEGATIVE • Units kJ mol-1

  45. Calculating ΔG • Method 1 use ΔGΘ = ΔHΘ – TΔSΘ • Method 2 use ΔGΘ = ΣΔGfΘ(products) – ΣΔGfΘ(reactant) • Standard free energy of formation the free energy change for the formation of 1 mol of substance from its elements in their standard state under standard conditions • Temp must be in K • If no temp is given and it is standard conditions, assume 298K

  46. Non-spontaneous rxns • If a rxn is non-spontaneous, does that mean it will never happen? NO • What will make it run? Outside influence such as; temp, catalysts, etc.

  47. Effects of temp on spontaneity • Refer to the equation: • ΔG = ΔH – TΔS • If ΔS is positive, temp must be high or low to be spontaneous? • High • ΔG must be negative to be spontaneous, thus TΔS must be higher than ΔH

  48. Temp and spontaneity • If ΔS is negative --TΔS will be positive and the rxn cannot be spontaneous • Endo rxns will only occur spontaneously in entropy is increased and temp is significantly high • Exorxns will always be spontaneous at some temp • If the rxn involves a decrease in entropy the rxn will be spontaneous at a lower temp; becomes less spontaneous as temp increases

  49. Signs

More Related