Reaction Predictions ! !

1. Reaction Predictions ! !

2. Types of Chemical Reactions • Single Displacement • Double Displacement • Decomposition • Synthesis • Combustion

3. Single Displacement • A +BC  AC + B • One element shoves the other element out!

4. Example 1: • Ca+AlCl3CaCl2+Al

5. Example 2: • AlCl3 + K

6. What do metals have to do with reaction predictions? • Metals vary in reactivity • Lose electrons in a reaction • AND give the electrons to someone else • Metal’s ability to give electrons depends on how reactive the metal is • How easily/quickly the metal wants to give up its electrons

7. Activity Series of Metals • Metals listed and arranged according to reactivity. • Metals will displace other metal ions in a solution from any metal BELOW it • Generally, • Metals > H2 on activity series will produce hydrogen gas (H2) when combined with an acid • Metals < H2 cannot produce hydrogen gas (H2) from an acid

8. Double Displacement • AB + CD  AC + BD • Elements switch partners !

9. Example 1: • Na3PO4+BaCl2Ba3(PO4)2 +NaCl

10. Example 2: K3PO4+ MgSO4

11. Decomposition • AB  A + B • Breaking chemical compound up, going from BIG to SMALL !

12. Example 1: • Au2O3Au + O2

13. Example 2: • H2O

14. Synthesis • A + B  AB • Joining! Making new chemical compound !

15. Example 1: • Mg + N2 Mg3N2

16. Example 2: • Be + Cl2

17. Combustion • HYDROCARBON (compound made up of just Cs and Hs) + O2CO2 + H2O • Chemical reactions involve a compound burning.

18. Example 1: • C2H6+ O2CO2+ H2O

19. Example 2: • Propane (C3H8) burns

20. Thermodynamics

21. Energy • Ability to do work • Units– Joules (J), we will use “kJ” • Can be converted to different types • Energy change results from forming and breaking chemical bonds in reactions

22. Heat (q) • Energy transfer between a system and the surroundings • Transfer is instant from high----low temperature until equilibrium • Temperature— • Measure of heat, “hot/cold” • the average kinetic energy of molecules

23. Heat (q) continued • Kinetic theory of heat • Heat increase resulting in temperature change causes an increase in the average motion of particles within the system. • Increase in heat results in • Energy transfer • Increase in both potential and kinetic energies

24. Thermodynamics 101 • First Law of Thermodynamics • Energy is conserved in a reaction (it cannot be created or destroyed)---sound familiar??? • Math representation: ΔEtotal = ΔEsys + ΔEsurr = 0 • Δ= “change in” • ΔΕ= positive (+), energy gained by system • ΔΕ= negative (-), energy lost by system • Total energy = sum of the energy of each part in a chemical reaction

25. Calorimetry How do we find the change in energy/heat transfer that occurs in chemical reactions???

26. Calorimetry • Experimentally “measuring” heat transfer for a chemical reaction or chemical compound • Calorimeter • Instrument used to determine the heat transfer of a chemical reaction • Determines how much energy is in food • Observing temperature change within water around a reaction container ** assume a closed system, isolated container • No matter, no heat/energy lost • Constant volume

27. Specific Heat • Amount of heat required to increase the temperature of 1g of a chemical substance by 1°C • Units: cal/g-K or J/g-K • 4.184 J = 1 cal, K = 273 + °C • Allows us to calculate how much heat is released or absorbed by a substance ! ! ! • Unique to each chemical substance • Al(s) = 0.901J/g°K • H2O(l) = 4.18 J/g°K

28. Specific Heat Equations • q = smΔΤ • s/Cp = specific heat (values found in reference table) • m = mass in grams • ΔΤ= change in temperature

29. Example 1: How much energy is required to warm 420 g of water in a water bottle from 25C to 37C ? Q = ? m = 420 g C(H2O (l)) = 4.18 J/g• C ΔT = 37-25 = 12 C Q = mc ΔT Q = (420 g)(4.18 J/g•  C)(12 C) Q = 21067 J or 21 kJ

30. “Coffee Cup” calorimeter • Styrofoam cup with known water mass in calorimeter • Assume no heat loss on walls • Initial water temp and then chemical placed inside • Final temperature recorded • Any temperature increase has to be from the heat lost by the substance SOOO • All the heat lost from the chemical reaction or substance is transferred to H2O in calorimeter

31. Example 3: • The specific heat of gold is 0.128 J/g°C. How much heat would be needed to warm 250.0 g of gold from 25°C to 100°C?

32. Heat of Fusion (Hf) • Fusion means melting/freezing • amount of energy needed to melt/freeze 1g of a substance • Different for every substance – look on reference tables • Q = mHf

33. Heat of Vaporization(Hv) • Vaporization means boiling/condensing • amount of energy needed to boil/condense 1g of a substance • Different for every substance – look on reference tables • Q = mHv

34. Examples: • Calculate the mass of water that can be frozen by releasing 49370 J. • Calculate the heat required to boil 8.65 g of alcohol (Hv = 855 J/g). • Calculate the heat needed to raise the temperature of 100. g of water from 25 C to 63 C .

35. Phase Change Diagram • The flat points represent a phase change – temperature does not change while a phase change is occurring even though heat is being added. • Diagonal points represent the 3 phases

36. Enthalpy Thermodynamics

37. Enthalpy (ΔH) • 2 types of chemical reactions: • Exothermic—heat released to the surroundings, getting rid of heat, -ΔΗ • Endothermic—heat absorbed from surroundings, bringing heat in, +ΔΗ **Enthalpy of reaction— amount of heat from a chemical reaction which is given off or absorbed, units = kJ/mol

38. Enthalpy of Reaction DH = Hfinal - HinitialHinitial = reactants Hfinal = products IfHfinal > Hinitial then DH is positive and the process is ENDOTHERMIC IfHfinal < Hinitial then DH is negative and the process is EXOTHERMIC

39. Enthalpy of Reaction Hfinal < Hinitial and DH is negative

40. Enthalpy of Reaction Hfinal > Hinitial and DH is positive

41. More Enthalpy • The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change • HgOHg + ½ O2 ΔH = + 90.83 kJ • Hg + ½ O2HgOΔH = - 90.83 kJ • SOOO-----total ΔH = 0

43. Methods for determining ΔH • Calorimetry • Application of Hess’ Law • Enthalpies of Formation

44. Guidelines for using Hess’ Law • Use data and combine each step to give total reaction • Chemical compounds not in the final reaction should cancel • Reactions CAN be reversed but remember to reverse the SIGN on ΔH

45. USING ENTHALPY Calculate DH for S(s) + 3/2O2(g)  SO3(g) knowing that S(s) + O2(g)  SO2(g) DH1 = -296.8 kJ SO2(g) + 1/2O2(g)  SO3(g) DH2 = -98.9 kJ The two equations add up to give the desired equation, so DHnet = DH1 + DH2 = -395.7 kJ

46. Example 4: NO(g) + ½ O2 NO2(g) ΔH° = ? Based on the following: ½ N2(g) + ½ O2 NO (g) ΔH° = + 90.29 kJ ½ N2(g) + O2NO2(g) ΔH° = +33.2 kJ

47. Enthalpy of Formation (ΔHf°) • Enthalpy for the reaction forming 1 mole of a chemical compound from its elements in a thermodynamically stable state. • A chemical compound is formed from its basic elements present at a standard state (25°C, 1 atm) • Enthalpy change for this reaction = ΔHf° • ΔHf°= 0 for ALL elements in their standard/stable state.

48. Enthalpy of Formation cont. • DHrxn = Hfinal – Hinitial Really, • ΔHf (products) - ΔHf(reactants) • Calculate ΔHrxn based on enthalpy of formation (ΔHf) • aA + bBcC + dD ΔH° =[c (ΔHf°)C + d(ΔHf°)D] - [a (ΔHf°)A+ b (ΔHf°)B ]

49. Entropy

50. Spontaneous vs. Nonspontaneous • Spontaneous Process • Occurs WITHOUT help outside of the system, natural • Many are exothermic—favors energy release to create an energy reduction after a chemical reaction • Ex. Rusting iron with O2 and H2O, cold coffee in a mug • Some are endothermic • Ex. Evaporation of water/boiling, NaCl dissolving in water