6.1 Energy Changes and Chemical Reactions

1. 6.1 Energy Changes and Chemical Reactions • All reactions release or absorb energy. • Common forms of energy are: heat, light, electricity, sound • Energy is stored in the bonds of compounds. • Breaking a bond requires energy. • Bonding releases energy.

2. The difference between the amount of energy absorbed in breaking the bonds in reactants and the amount of energy released in forming bonds determines whether the reaction is exothermic or endothermic. • Reactions which need more energy to break bonds than they make when bonds are formed are endothermic (require heat). • Reactions that make more energy as bonds form than is needed to break the bonds in the reactants are exothermic (give off heat).

4. Consider the following: • It is easier (requires less energy) to build small stable compounds. • It is harder (requires more energy) to build large unstable compounds. • Explosives are typically large, unstable compounds that form small stable compounds when they react. • The leftover energy drives the explosion.

5. TNT (Trinitrotoluene) C6H2(CH3)(NO3)2(s) + O2(g) → CO2(g) + H2O(g) + N2(g) Very complex, unstable compoundsimple, stable compounds • Lots of E required to break the reactant bond • Little energy required to make product bonds • Leftover E drives the explosion

6. Law of Conservation of Energy Energy cannot be created or destroyed. It can only be changed from one form to another. Example: Burning Wood ReactantsProducts Chemical energy in wood heat energy chemical energy radiant energy sound energy

7. Magic Wand Demonstration • Read pages 180 - 182 • Assignment: • p 185 #1 – 5d & BLM do sheet 6-1, read 6-2 and do 6-4 • 6-3 may be done as a bonus assignment

8. 6.2 Synthesis and Decomposition Reactions • Synthesis Reaction – synthesis means to make or blend – these chemical reactions have several small reactants combining to form larger products. General Form: If A is an element, and X is an element we would have: A + X AX

9. Example: When pop is made: CO2(g) + H2O(l) H2CO3(aq)

10. Decomposition Reaction • Decompose means to break down. • These chemical reactions begin with larger reactant(s) which break down into simpler products.

11. General From: • AX A + X Example: • When you open up the pop can: • H2CO3(aq) CO2(g) + H2O(l) • The CO2 coming out of solution produces the fizz.

12. NOTICE the brackets give the state of the substance. • (s) is solid • (l) is liquid • (g) is gas • (aq) is aqueous which means dissolved.

13. Recall: • A solid takes up little volume, • The same mass of liquid occupies more volume. • A same mass of a gas occupies much more volume. • Recall: • Materials expand as they get hot. • Explosions are biggest when the reaction can change a material of small volume to a material with large volume VERY QUICKLY. • Rockets, firearms, explosives, etc.

14. Assignment • Read p 186 – 188 • Assignment p 189 #1 – 5 • BLM sheet 6-5 & 6-6

15. 6-3 Single and Double Displacement Reactions • a chemical reaction in which one element takes the place of another element in a compound. • Example: Let A and B be cations. Let X be an anion. • A + BX  AX + B

16. Think A & B are guys. X is a girl. B and X are dance partners and have “bonded”. BUT along comes A. Looks like A and X bond and B is left alone. • How could we predict if this would happen or not? Well A would have to be more “reactive” with X than B was. • In chemistry the reactivity of elements has been tested and the results are available in reactivity charts called The Activity Series of Metals.

17. Real Life Example: • Mg is more reactive than hydrogen so: • Mg(s) + HCl(aq)  MgCl2(aq)+ H2(g) • Balance the above reaction. • 1Mg(s) + 2HCl(aq) 1MgCl2(aq)+ 1H2(g)

18. Example 2: • Na is very reactive so: • Na(s) + H2O(l)  NaOH(aq) + H2(g) • Think of water as being ionic Hydrogen Hydroxide: • Na(s) + H(OH)(l)  Na(OH)(aq) + H2(g) • Balance the above reaction. • 2Na(s) + 2H(OH)(l) 2Na(OH)(aq) + 1H2(g)

19. Double Displacement Reaction – a reaction in which the cations from 2 different compounds exchange forming 2 new compounds. • AX + BY  BX + AY • Think of 2 couples and the guys switched places. • AX + BY

20. Example: • 2KI + 1Pb(II)(NO3)21Pb(II)I2 + 2K(NO3) • Notice: an interesting event takes place in the reaction. Two liquids are mixed and one of the materials produced is a solid. The solid forms as particles that slowly “rain” down to the bottom of the liquid. The solid is called a precipitate.

21. Neutralization • Neutralization – a double displacement reaction that occurs when an acid and a base are combined. The products are always the same: • Acid + Base  Water + Salt • A salt is an ionic compound.

22. Ex) • HNO3 + NaOH  H2O + NaNO3 • (NOTE: it often helps to think of H2O as H+OH-) • Read P 190 – 202 and do P 202 1-4, BLM 6-7, 6-8 (read), 6-9 & 6-12

23. 6.4 Organic Chemistry • Organic molecules contain carbon (carbon is found in all living things). • Carbon Bonding • Carbon has 4 valence electrons . .C. .

24. Carbon does not form ions, but bonds through covalent bonds (shares e-) • Carbon forms chains with other carbon atoms • Examples: Methane - CH4 H lX H lXC lX H lX H

25. Ethane – C2H6 • Draw the Lewis dot diagram for this • We will use carbon chain drawings • See board • Carbon chains are most often found with hydrogen

26. Atoms complete all “open ends” to the chains • Examples: butane (C4H10) and propane (C3H8) • Draw these • Hydrocarbons (CH) are carbon chains that contain hydrogen • Hydrocarbons are very useful fuels

27. Combustion • Combustion means burning • Hydrocarbons are excellent fuels • When fuels burn they combine with oxygen in exothermic reactions • HC + O2 CO2 + H2O + Thermal E • We use the thermal energy: • for heating • to expand gas quickly and we use the force produced by the expanding gas (internal combustion engine)

28. Complete Combustion • In a perfect world every hydrocarbon molecule would combine with oxygen to produce the maximum amount of heat, and only CO2 and H2O

29. Incomplete Combustion • Reactions are often limited by a limited amount of one of the reactants (O2 or fuel). • Incomplete combustion: • reduces thermal energy • wastes fuel (fuel in vehicle exhaust) • produces Carbon monoxide (CO) which is deadly – odorlous, colourless, binds with rec blood cells better than O2 but is not useable by the body) • produces carbon (soot) and causes smog and respiratory diseases including cancer

30. HC + O2 CO2 + H2O + CO + C (soot) + Reduced Thermal E

31. Assignment • CYU P 206 1 – 4 & BLM 6-14, 6-15, 6-16 Ch 6 Review • BLM 6-21 • Ch Review P 209 1-18 • Pop Can Canon Demonstration

32. Ch 6 - Lab Copper (II) Chloride – CuCl2 Sodium Hydroxide – NaOH Sodium Chloride – NaCl Potassium Nitrate – KNO3 Sodium Carbonate – Na2CO3 Calcium Chloride – CaCl2 Potassium Iodide – KI Lead (II) Nitrate – Pb(NO3)2

33. H2 and O2 Gas Test • Wooden splint – glows/ignites (2) • Lighter – explodes/”pops” (2) • Mg and O2 • Exothermic, flame and light, heat (2) • 2Mg + O2 2MgO + thermal E (4) • Ionic, metal/non-metal (2) • 3Mg + N2 = Mg3N2 + thermal E (4)

34. Mg and HCl • Bubbles and heat (2) • Single displacement, Mg and H switch places (2) • Mg and HCl (1) • MgCl2 and H2 (1) • Mg + 2HCl = MgCl2 + H2 + thermal E (4)

35. Dble. Displacement • Blue precipitate (gelatin) (2) • CuCl2 + 2NaOH = Cu(OH)2 + 2NaCl (4) • None (2) • NaCl + KNO3 = NR (4) • Turned foggy/milky, white precipitate (2) • Na2CO3 + CaCl2 = CaCO3 + NaCl (4) • Yellow precipitate (2) • 2KI + Pb(NO3)2 = PbI2 + 2KNO3 (4)