Combustion Reactions (another example of redox) Combustion

Transcription

1 Combustion Reactions (another example of redox) Combustion is the burning of a fuel by rapid oxidation with oxygen in air. - General reaction for carbon based fuels is: Balancing Combustion Reactions Step 1: Balance C Step 2: Balance H Step 3: Balance O Step 4: Change any half numbers to a whole number and check to be sure of balancing. Examples: C 7 H 16 (l) + O 2 (g) CO 2 (g) + H 2 O(g) C 4 H 10 (g) + O 2 (g) CO 2 (g) + H 2 O(g) C 4 H 10 (g) + O 2 (g) CO 2 (g) + H 2 O(g) Problem: balance the following oxidation reaction: C 2 H 6 (g) + O 2 (g) CO 2 (g) + H 2 O(g) Ch 5 Page 15

2 5.6 Recognizing Redox Reactions How can you tell if a reaction involves a transfer of electrons? - When ions are involved, determine if there is a change in. - All single replacement reactions are redox reactions. - Single replacement reactions exchange which type of atom is present as a pure element. An example of a single replacement reaction: - For molecular substances, a value called an (or oxidation state), indicates whether the atom is neutral, electron-rich, or electron-poor. - By comparing the oxidation number of an atom before and after a reaction, we can tell whether the atom has gained or lost shares in electrons. - Oxidation numbers do imply ionic charges. hey are simply a convenient device for keeping track of electrons in redox reactions. Rules for Assigning Oxidation umbers 1. Elements = 0 e.g. K, 2, Xe (oxid # = 0) 2. All oxidation numbers add to zero (or to ion charge for an ion). 3. A fixed charge cation retains its charge. Group #1 (oxid # = +1) Group #2 (oxid. # = +2) Ag ion = +1; Cd ion = +2; Zn ion = +2; Al ion = Hydrogen attached to a nonmetal (oxid # = +1) Hydrogen attached to a metal (oxid # = -1) 5. Oxygen (usually has oxidation # = -2) (unless it violates #2; e.g. peroxides) 6. Halide (F always -1; others usually has oxidation # = -1) (unless it violates #2) 7. Determine remaining element oxidation # by following rule #2. Ch 5 Page 16

3 Examples aclo4 aclo3 aclo2 Molecular substances can be analyzed in terms of loss and gain of oxygen. - he oxygens in the formula, the the oxidation state, so the more oxidized the substance. Problen: Determine the oxidation number of V in VCl 3. Problem Determine the oxidation number of phosphorus in a 3 PO 4. Problem: Determine the oxidation number of in O 3 -. Problem: Determine the oxidation number of Cr in Cr 2 O Remember that the oxidation number is the number for a single atom! Ch 5 Page 17

4 Using oxidation numbers for redox recognition. dentify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent in the following reaction. Cu + O 3 - Cu 2+ + O 2 oxidized: reducing agent: reduced: oxidizing agent: Agents include the entire solids, covalent molecules & polyatomic ions Problem: dentify the oxidizing agent and the reducing agent in the following reaction. Fe 2 O 3 (s) + 3 CO(g) 2 Fe(s) + 3 CO 2 (g) is the oxidizing agent; is the reducing agent. he Activity Series Redox as a Battle for Electrons We have learned that total molecular equations can be broken down. An equation such as: 2 Al + 3 CuCl 2 2 AlCl Cu Can be written as a net ionic equation: Ch 5 Page 18

5 2 Al + 3 Cu 2+ à 2 Al Cu Redox reactions can be broken down even further, to include separate oxidation and reduction half reactions. Oxidation ½ Rxn: Reduction ½ Rxn: How do we know the reaction goes in this direction? Why not write it as: 2 Al Cu à 2 Al + 3 Cu 2+ Problem: try writing the half reactions for the reverse reaction. Oxidation ½ Rxn: Reduction ½ Rxn: his is just as valid on paper. Let s see what reaction actually occurs! DEMO: Fill 1 beaker full of CuCl 2 (source of Cu 2+ ions) and add some Al foil. (his represents scenario 1 above) Reaction? Fill beaker 2 with Al(O 3 ) 3 (source of Al 3+ ions) and add some Cu metal. Cu metal will quickly form on the surface of the Al in beaker 1. Reaction? his shows us that is a more reactive metal than. is more easily oxidized. (t wants to form a cation the most.) n order to put elements into order of the greatest activity, we put the ones that want to be cations the most at the top of the list. We will practice doing this during lab: (Single and Double Replacement Reactions.) Ch 5 Page 19

6 Driving Forces f a reaction occurred, then there must have been something that made it more than just a mixture when reactants are added together. hat something is known as a driving force. here are four general driving forces. 1. (ons stick together and fall to the bottom of the beaker.) 2. (ons stick together and leave through the top of the beaker as bubbles. 3. (ons come together as a covalent molecule and stick together while floating around) 4. (Electrons are exchanged) Did a Reaction Occur? n order for a reaction to occur, the reactants and the products must be exactly the same. f we have all the same species floating around in solution before and after a potential reaction, then no reaction actually occurred! Let s learn how to write equations that leave out non-reacting spectators. 5.7 et onic Equations n reactions involving ions, it is more accurate to write the reaction as an. An ionic equation is one in which ions are explicitly shown as entities. Some ions undergo no change during the reaction. hey appear on sides of the reaction because they play no role. hese are known as ions. he actual reaction can be described more simply by writing a net ionic equation, which spectator ions. Ch 5 Page 20

7 A net ionic equation must be balanced both for and for, with all coefficients reduced to their lowest whole numbers. hings that break apart in solution - break into H + and A - Strong acids are HCl, HBr, H, HClO 4, HClO 3, HO 3, H 2 SO 4 - break into M + and OH - Common strong bases are aoh, KOH, LiOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2, - break up into their ions. Use solubility rules to decide hings that stay together in solution - Only aqueous species break up! - Formulas that start with H that are not listed above. - Formulas ending in OH not listed above. - Water is the most common one we will find. - Any substance that is insoluble and would form a precipitate. When spectator ions are removed from the equation we are left with a. Ch 5 Page 21

8 SEPS O WRG E OC EQUAOS 1. Convert the reactants from words to chemical formulas. 2. Determine the type of reaction that is likely to occur and use it to predict the products. 3. Balance the reaction. - his is often called the total or molecular equation. 4. Use your solubility rules to determine the appropriate state symbols. - insoluble substances will have (s). - soluble salts will have (aq). 5. Break into anions and cations those things that are primarily broken apart in solution. - Break apart soluble salts, strong acids, and strong bases. - Don t break apart insoluble substances, weak acids, weak bases, gases, and H 2 O. Be sure to take stoichiometric coefficients into account 2 a 2 SO 4 becomes 4a + (aq) + 2SO 4 2- (aq) - his is called the otal onic Equation. 6. Simplify the equation by cancelling out things that appear on both sides of the reaction. - hey must appear with the same charge and in the same phase to cancel out! - tems that cancel out are considered spectators since they don t actively participate. - he resulting reaction is the et onic Equation. Ch 5 Page 22

9 For example, let look at the reaction of ammonium iodide with lead () perchlorate: --otal Equation: --onic Equation: --et onic Equation: ectators Problem: Write net ionic equations for the following 4 equations. nclude state symbols for the total equation and state the driving force. a) Solutions of potassium hydroxide and aluminum nitrate were mixed together. b) Solutions of phosphoric acid and calcium hydroxide neutralize each other. Ch 5 Page 23

10 c) Solutions of sodium chloride and potassium nitrate are mixed. d) Solutions of calcium acetate and lithium sulfide react. e) Elemental copper reacts with a solution of silver nitrate. Problem: Determine the net ionic equation for: a 2 SO 4 (aq) + Ba(O 3 ) 2 (aq) Ch 5 Page 24