1 / 36

Methane

Methane. Hydrocarbons – compounds containing only carbon and hydrogen. hydrocarbons. aromatic. aliphatic. alkanes. alkenes. alkynes. Alkanes – hydrocarbons with the general formula C n H 2n+2 (four bonds to each carbon and only single bonds) CH 4 methane C 2 H 6 ethane

eshe
Download Presentation

Methane

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. Methane

  2. Hydrocarbons – compounds containing only carbon and hydrogen. hydrocarbons aromatic aliphatic alkanes alkenes alkynes

  3. Alkanes – hydrocarbons with the general formula CnH2n+2 (four bonds to each carbon and only single bonds) CH4 methane C2H6 ethane C3H8 propane Etc.

  4. Methane = CH4 H | H—C—H sp3 tetrahedral 109.5o bond angles | H Non-polar – van der Waals (London forces) Gas at room temperature mp = -183oC bp = -161.5oC Water insoluble Colorless and odorless gas “swamp gas” ; fossil fuel found with petroleum & coal Important fuel/organic raw material

  5. Chemistry of methane (reactions)? CH4 + H2O  CH4 + conc. H2SO4 CH4 + conc. NaOH  CH4 + sodium metal  CH4 + KMnO4  CH4 + H2/Ni  CH4 + Cl2  NR (no reaction) NR NR NR NR NR NR

  6. Methane is typically unreactive. It does not react with water, acids, bases, active metals, oxidizing agents, reducing agents, or halogens. • Reactions of methane: • Combustion (oxidation;complete & partial) • Halogenation

  7. Reactions of Methane • Combustion (oxidation) • a) complete oxidation • CH4 + 2 O2 , flame or spark  CO2 + H2O + energy • b) partial oxidation • 6 CH4 + O2 , 1500o  CO + H2 + H2C2 (acetylene) • CH4 + H2O , Ni, 850o  CO + H2

  8. Halogenation • CH4 + X2 , Δ or hυ  CH3X + HX • X2 = Cl2 or Br2 •  a) Requires heat (Δ) or uv light (hυ) • b) May proceed further •  c) Cl2 reacts faster than Br2 •  d) No reaction with I2

  9. “Substitution” reaction CH4 + Cl2 CH4 + I2, heat  CH4 + Br2, hv  NR (requires heat or uv light) NR (does not react with I2) CH3Br + HBr

  10. CH4 + Cl2, hv  CH3Cl + HCl methyl chloride chloromethane CH3Cl + Cl2, hv  CH2Cl2 + HCl methylene chloride dichloromethane CH2Cl2 + Cl2, hv  CCl3H + HCl chloroform trichloromethane CCl3H + Cl2, hv  CCl4 + HCl carbon tetrachloride tetrachloromethane

  11. CH4 + Br2, hv  CH3Br + HBr methyl bromide bromomethane CH3Br + Br2, hv  CH2Br2 + HBr methylene bromide dibromomethane CH2Br2 + Br2, hv  CBr3H + HBr bromoform tribromomethane CBr3H + Br2, hv  CBr4 + HBr carbon tetrabromide tetrabromomethane

  12. CH3I CH2I2 iodomethane diiodomethane methyl iodide methylene iodide CHI3 CI4 triiodomethane tetraiodomethane iodoform carbon tetraiodide

  13. Can proceed further: CH4 + Cl2, heat  CH3Cl + CH2Cl2+ CHCl3+ CCl4 + HCl Control? (xs) CH4 + Cl2, heat  CH3Cl + HCl bp –162o bp –24o CH4 + (xs) Cl2, heat  CCl4 + 4 HCl

  14. Mechanism for the monochlorination of methane initiating step: • Cl2 2 Cl• propagating steps: • Cl• + CH4 HCl + CH3• • CH3• + Cl2 CH3Cl + Cl• then 2), then 3), then 2), etc. terminating steps: • Cl• + Cl•  Cl2 • Cl• + CH3•  CH3Cl • CH3• + CH3•  CH3CH3

  15. Energy Changes? ΔH Homolytic bond dissociation energies (see inside the front cover of M&B) H—Cl 103 Kcal/mole Cl—Cl 58 Kcal/mole CH3—H 104 Kcal/mole CH3—Cl 84 Kcal/mole

  16. We need only consider those bonds that are broken or formed in the reaction. CH3—H + Cl—Cl  CH3—Cl + H—Cl +104 +58 -84 -103 PE: +162 -187 ΔH = +162 –187 = -25 Kcal/mole (exothermic, gives off heat energy)

  17. ΔH for each step in the mechanism? • Cl—Cl  2 Cl• +58 ΔH = +58 • Cl• + CH3—H  H—Cl + CH3• +104 -103 ΔH = +1 • CH3• + Cl—Cl  CH3—Cl + Cl• +58 -84 ΔH = -26 • Cl• + Cl•  Cl—Cl -58 ΔH = -58

  18. Rates of chemical reactions depend on three factors: Collision frequency (collision per unit time) Probability factor (fraction of collisions with correct geometry) Energy factor (fraction of collisions with sufficient energy) “sufficient energy” = Energy of activation, minimum energy required for a collision to go to the product.

  19. Z = collision frequency P = probability factor e-Eact/RT = fraction of collisions with E > Eact Note: rate decreases exponentially as the Eact increases!

  20. @ 275oC Eact Collisions > Eact 5 Kcal 10,000/1,000,000 10 Kcal 100/1,000,000 15 Kcal 1/1,000,000 If the Eact is doubled, the rate is decreased by a factor of 100 times!

  21. Eact cannot be easily calculated like ΔH, but we can estimate a minimum value for Eact: If ΔH > 0, then Eact > ΔH If ΔH < 0, then Eact > 0

  22. Rate determining step (RDS) = the step in the mechanism that determines the overall rate of a reaction. In a “chain reaction” this will be the slowest propagating step. For chlorination of methane, which propagating step is slower? Step 2) ΔH = +1 Kcal/mole Eact > +1 Kcal (estimated) Step 3) ΔH = -26 Kcal/mole Eact > 0 Kcal (estimated) Step 2 is estimated to be slower than step 3 and is the RDS

  23. An “alternate mechanism: • Cl• + CH4 CH3Cl + H• • H• + Cl2 HCl + Cl• • Why not this mechanism? • Step 2: ΔH = +104-84 = +20 Kcal/mole; Eact > +20 Kcal • Step 3: ΔH = +58-103 = -45 Kcal/mole; Eact > 0 Kcal • RDS for this mechanism is step 2 and requires a minimum of 20Kcal/mole! Unlikely compared to our mechanism where the RDS only requires an estimated minimum of 1 Kcal!

  24. Halogenation • Δ or hυ • CH4 + X2 CH3X + HX • requires heat or light • X2: Cl2 > Br2 I2 • why?…how?…mechanism

  25. This reaction requires heat or light because the first step in the mechanism involves the breaking of the X-X bond. This bond has to be broken to initiate the chain mechanism. F—F 38 Kcal/mole Cl—Cl 58 Kcal/mole Br—Br 46 Kcal/mole I—I 36 Kcal/mole Once initiated the reaction may or may not continue based on the Eact for the RDS.

  26. “generic” mechanism for the halogenation of methane • (free radical substitution mechanism) • X2 2 X• • X • + CH4 HX + CH3• • CH3• + X2  CH3X + X• • 2 X•  X2 • X• + CH3•  CH3X • 2 CH3•  CH3CH3

  27. ΔH for each step in the mechanism by halogen: F Cl Br I 1 +38 +58 +46 +36 2 -32 +1 +16 +33 3 -70 -26 -24 -20 4 -38 -58 -46 -36 5 -108 -84 -70 -56 6 -88 -88 -88 -88

  28. Estimation of Eact for the propagating steps: Eact (est.) F Cl Br I 2 >0 >+1 >+16 >+33 3 >0 >0 >0 >0 Step 2 is the RDS Rate Cl2 > Br2 because in the RDS Eact(Cl2) < Eact(Br2) NR with I2 because RDS Eact(I2) > +33 Kcal/mole only 1/1012 collisions would have E > +33 at 275o

  29. The transition state (‡) or “activated complex” is the unstable structure that is formed between reactants and products in a step in a mechanism. It corresponds to the energy at the top of the energy barrier between reactants and products. step 2 in the chlorination of methane: Cl• + CH4 HCl + CH3• Transition state: [ Cl--------H-------CH3 ]‡ δ• δ•

  30. Hammond’s Postulate: the higher the Eact of a step in a mechanism, the later the transition state is reached and the more the transition state will look like the products. In step 2 of the mechanism for the bromination of methane, the Eact is estimated to be > +16 Kcal/mole. Since the Eact is high, the transition state is reached later in this step than it is in chlorination and will look more like the products: [ Br----H-----------CH3 ]‡ δ•δ•

  31. Reactions of Methane • Combustion (oxidation) • a) complete oxidation • CH4 + 2 O2 , flame or spark  CO2 + H2O + heat • b) partial oxidation • 6 CH4 + O2, 1500oC  CO + H2 + H2C2 • CH4 + H2O, 850o, Ni  CO + H2 • Halogenation • CH4 + X2, heat or hv  CH3X + HX • requires heat or light • Cl2 > Br2 NR with I2

More Related