1 / 23

CHEMICAL REACTIONS

Learn about chemical reactions, bond energy, thermodynamics, and equilibrium in this comprehensive guide. Explore factors affecting reaction rates, equilibrium position, and the significance of equilibrium constants.

bkarla
Télécharger la présentation

CHEMICAL REACTIONS

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. CHEMICAL REACTIONS • When chemical reactions occurOLDbonds (in thereactants) are broken andNEWbonds (in theproducts) are formed. • The energy needed to break old bonds and form new ones can be studied throughTHERMOCHEMISTRY.

  2. Bonds • Ionic - electrostatic forces of attraction between ions • Covalent - forces of mutual attraction of electrons between NONMETAL atom s (most NM atoms want an OCTET) • Hydrogen - weak forces of attraction • between water molecules • within DNA, holding the 2 strands together

  3. Energy • Potential Energy (stored energy) - the energy of Chemical Bonds • Bond formation always releases energy; exothermic • Bond dissociation always requires energy; endothermic • Kinetic Energy (energy associated with motion): KE = 1/2mv2

  4. THERMOCHEMISTRY • The study of heat changes during chemical reactions. • Based on the net energy of bonds dissociating and bonds breaking • ∆H is the symbol representing “change in heat” • Differences in bond dissociation energies allow us to determine if heat is/will be released or needs to be absorbed during a reaction. • ∆H = B.E.PB.E.R • If the products require less bond energy than the reactants, the excess energy is released (∆H = -) and vice versa (∆H = +)

  5. Endothermic Reactions • The reacting chemicals absorb heat from their surroundings (Heat In!) • ∆H = + • Ba(OH)2 + NH4Cl + heat --> NH3 + BaCl2 + H2O

  6. Exothermic Reactions • The reacting chemicals release heat into their surroundings (Heat Out!) • ∆H = - • KMnO4 + C3H8O3 --> K2CO3 + Mn2O3 + CO2 + H2O + heat • C12H22O11H2SO4 > C + H2O

  7. Energy graphs showing the difference between an exothermic and an endothermic reaction. KMnO4 + C3H8O3 --> K2CO3 + Mn2O3 + CO2 + H2O + heat Ba(OH)2 + NH4Cl + heat --> NH3 + BaCl2 + H2O

  8. Rate of Reaction • Rate = Speed • The rate of a reaction depends on: • Temperature, Concentration of reactants, Catalysts • Reactions require a specific amount of “activation” energy (Ea) in order for reactants to react effectively.

  9. Rubbing a match head against a rough surface provides the activation energy needed for the match to ignite.

  10. Factors that Affect Rxn Rates

  11. Rxn Rates & Concentration Graphs showing how reaction rates and reactant concentration vary with time.

  12. Catalysts Catalysts lower the activation energy for chemical reactions.

  13. Equilibrium • Many chemical reactions occur in two directions - forward and reverse. • Once the reaction is established an equilibrium can develop. • Rate of forward reaction = Rate of reverse reaction A(aq) + B(aq) <==> AB(aq)

  14. 3 factors affect equilibrium • Concentration (substances must be in aqueous or gaseous form) • Temperature (exo vs. endo) • Pressure - affects gases only (look at the # of moles of gases) A(aq) + B(aq) <==> AB(aq)+ heat

  15. Effect of Concentration Changes Concentration changes that result when H2 is added to an equilibrium mixture.

  16. Effect of Temperature Equilibrium mixtures changing color with difference in temperatures.

  17. Equilibrium Position • A + 2B <==> C + D • This position is defined by the amounts of reactants and products • If the equilibrium position shifts, equilibrium will have to be reestablished with different amounts of reactants and products. • An equilibrium expression allows for a mathematical description of the position at equilibrium.

  18. Equilibrium Expression • A ratio of [products] over [reactants] • Each [ ] is raised to the power equal to its coefficient in the balanced equation • The ratio is set equal to a constant (Keq) A2(aq) + 2 B(aq) <==> 2AB(aq) Keq = Ex.: 2NOCl(g) <==> 2NO(g) + Cl2(g) BaCl2(aq) + Na2SO4(aq) <==> 2NaCl(aq) + BaSO4(s)

  19. Significance of Keq • 2NOCl(g) <==> 2NO(g) + Cl2(g) • If Keq = 1000, then the amount of products is essentially 1000x greater than reactants.

  20. Calculating Keq A2 + B2 <==> 2AB • Calculate Keq when [A] = 0.25 M; [B] = 0.35 M; [AB] = 2.50 M • If [A] increases to 0.55 M; [B] increases to 0.45 M, what would the new [AB] become?

  21. Chemical stress effects • Le Chatelier’s Principle: A system in equilibrium which is stressed tries to return to equilibrium by shifting the reaction in a direction to relieve the stress • So, if we increase the concentration of some participant in the equilibrium, the system will try to react away that substance. • If we decrease the concentration of some participant in the equilibrium, the system will try to produce more of that substance. • If we increase the temperature or pressure of the system, the system will try to reduce the temperature or pressure.

  22. Example of Le Chatelier’s Principle • C6H6(g) + 3H2(g) <==> C6H12(g) + heat • Increase [C6H6] • Decrease [C6H12] • Increase temperature

  23. Reactions of Ionic Compounds (an important example) • Tooth Enamel Demineralization Ca10(PO4)6(OH)2 <==> 10Ca2+ + 6PO43- + 2OH-

More Related