1. Chapter 6 Chemical Reactions Chemical Reactions

2. Chemical Reactions In a chemical reaction, one or more reactants is converted to one or more products In this chapter we discuss three aspects of chemical reactions (a) mass relationships (stoichiometry) (b) types of reactions (c) heat gain and loss accompanying reactions

3. Chemical Equations The following chemical equation tells us that propane gas and oxygen gas react to form carbon dioxide gas and water vapor But while it tells us what the reactants and products are and the physical state of each, it is incomplete because it is not balanced

4. Balancing Equations To balance a chemical equation – begin with atoms that appear only in one compound on the left and one on the right; in this case, begin with carbon (C) which occurs in C 3 H 8 and CO 2 – now balance hydrogens, which occur in C 3 H 8 and H 2 O – if an atom occurs as a free element, as for example Mg or O 2, balance this element last; in this case O 2

5. Balancing Equations Practice problems: balance these equations

6. Balancing Equations Solutions to practice problems – it is common practice to use only whole numbers; therefore, multiply all coefficients by 2, which gives

7. WILL A REACTION OCCUR? When zinc metal is added to hydrochloric acid, a reaction occurs: Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g) When copper metal placed in hydrochloric acid, no reaction occurs: Cu(s) + 2HCl(aq)  CuCl 2 (aq) + H 2 (g)

8. Oxidation-Reduction Who wants the electrons more? Oxidation: Oxidation: the loss of electrons Reduction: Reduction: the gain of electrons Oxidation-reduction (redox) reaction: Oxidation-reduction (redox) reaction: any reaction in which electrons are transferred from one species to another

9. Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a solution of copper(II) sulfate – some of the zinc metal dissolves – some of the copper ions deposit on the zinc metal – the blue color of Cu 2+ ions gradually disappears In this oxidation-reduction reaction – zinc metal loses electrons to copper ions – copper ions gain electrons from the zinc

10. Oxidation-Reduction – we summarize these oxidation-reduction relationships in this way

11. Oxidation-Reduction Although the definitions of oxidation (loss of electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others – for example, the combustion of methane An alternative definition of oxidation-reduction is – oxidation: – oxidation: the gain of oxygen or loss of hydrogen – reduction: – reduction: the loss of oxygen or gain of hydrogen

12. Oxidation-Reduction – using these alternative definitions for the combustion of methane

13. CLASSIFICATION OF CHEMICAL REACTIONS Oxidation Reduction Reactions Combination: N 2 + 3H 2  2NH 3 Decomposition: 2H 2 O  2H 2 + O 2 Substitution/displacement: Fe(s) + CuCl 2 (aq)  Cu(s) + FeCl 2 (aq) Double replacement (metathesis): Neutralization: HCl + NaOH  H 2 O + NaCl Precipitation: CaCl 2 (aq) + Na 2 CO 3 (aq)  CaCO 3 (s) + 2NaCl(aq) Evolution of gas: Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g )

14. Reactions Between Ions Ionic compounds, also called salts, consist of both positive and negative ions When an ionic compound dissolves in water, it dissociates to aqueous ions What happens when we mix aqueous solutions of two different ionic compounds? – if two of the ions combine to form a water-insoluble compound, a precipitate will form – otherwise no physical change will be observed

15. Reactions Between Ions Example: – suppose we prepare these two aqueous solutions – if we then mix the two solutions, we have four ions present; of these, Ag + and Cl - react to form AgCl(s) which precipitates

16. Reactions Between Ions – we can simplify the equation for the formation of AgCl by omitting all ions that do not participate in the reaction net ionic equation – the simplified equation is called a net ionic equation; it shows only the ions that react spectator ions – ions that do not participate in a reaction are called spectator ions

17. Reactions Between Ions In general, ions in solution react with each other when one of the following can happen – two of them form a compound that is insoluble in water – two of them react to form a gas that escapes from the reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid – an acid neutralizes a base (Chapter 8) – one of the ions can oxidize another (Section 4.7)

18. Reactions Between Ions Following are some generalizations about which ionic solids are soluble in water and which are insoluble – all compounds containing Na +, K +, and NH 4 + are soluble in water – all nitrates (NO 3 - ) and acetates (CH 3 COO - ) are soluble in water – most chlorides (Cl - ) and sulfates (SO 4 2- ) are soluble; exceptions are AgCl, BaSO 4, and PbSO 4 – most carbonates (CO 3 2- ), phosphates (PO 4 3- ), sulfides (S 2- ), and hydroxides (OH - ) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH 4 OH which are soluble in water

19. Energy The Capacity to do Work

20. The Two Types of Energy Potential: due to position or composition – One must compare what can be made from what you have in order to determine the potential. Kinetic: due to motion of the object KE = 1 / 2 mv 2 (m = mass, v = velocity)

21. Food Calories Food is stored chemical potential energy. Energy is produced by reacting the food we eat with oxygen. Some of the energy can be used to do work!

22. Exothermic Reactions Some reactions give off heat. Once the magnesium begins to burn, the heat given off makes more magnesium burn. Link to Magnesium burning

23. Endothermic Reactions Some reactions use heat during a reaction. The hot water supplies the heat needed to sublime the solid carbon dioxide into gaseous carbon dioxide.

24. Energy Potential - Stored energy – It is not the energy in the bonds of the compound you have in your hands that makes it have potential energy. – It is the energy of the new bonds that are formed in a chemical reaction that makes for the potential energy in the reactants.

25. Energy Diagrams It is the difference which determines the potential

26. Covalent Bond Strength Most simply, the strength of a bond is measured by determining how much energy is required to break the bond. This is the bond enthalpy. The bond enthalpy for a Cl—Cl bond is measured to be 242 kJ/mol.

27. The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. H 2 (g) H (g) +  H 0 = 436.4 kJ Cl 2 (g) Cl (g) +  H 0 = 242.7 kJ HCl (g) H (g) +Cl (g)  H 0 = 431.9 kJ O 2 (g) O (g) +  H 0 = 498.7 kJ OO N 2 (g) N (g) +  H 0 = 941.4 kJ N N Bond Energy Bond Energies Single bond < Double bond < Triple bond 9.10

28. Average Bond Enthalpies This table lists the average bond enthalpies (kJ/mol) for many different types of bonds.

29. Energy The energy required to break a bond is equal to the energy liberated when making a bond. LAW of conservation of energy – – Neither created nor destroyed during physical or chemical changes

30. Enthalpies of Reaction CH 4 (g) + Cl 2 (g)  CH 3 Cl (g) + HCl (g) In this example, One C—H bond and one Cl—Cl bond are broken; one C—Cl and one H—Cl bond are formed. So,  H rxn = [D(C—H) + D(Cl—Cl)  [D(C—Cl) + D(H—Cl) = [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)] = (655 kJ)  (759 kJ) =  104 kJ

31. Kinetics How Fast Will the Reaction Go?

32. HOW FAST DOES A REACTION GO? The reaction for the oxidation of sucrose, C 12 H 22 O 11 + 12O 2  12CO 2 + 11H 2 O does not occur at a perceptible rate at room temperature The reaction can take pace if you heat the sucrose with a flame to temperatures above 200 ˚C. The reaction also takes place at 37˚C through the action of enzymes in the body What effects the rate of chemical reactions?

33. Reaction Rates The rates of chemical reactions are affected by the following factors –molecular collisions –activation energy –nature of the reactants –concentration of the reactants –temperature –presence of a catalyst On the following screens, we examine these factors one at a time

34. Molecular Collisions –two species must collide to react –calculations show that the rate things collide is far greater than the rate at which they react –Conclusion: most collisions do not result in a reaction effective collision –a collision that results in a reaction is called an effective collision –there are two main reasons why some collisions are effective and others are not; energy orientation

35. Molecular Collisions Activation energy: Activation energy: the minimum energy required for a reaction to take place –energy is required for reactions to begin even if they give off energy during the process –this energy comes from collisions –if the collision energy is large, there is sufficient energy to break the necessary bonds, and reaction takes place –if the collision energy is too small, no reaction occurs

36. Molecular Collisions Orientation at the time of collision –the colliding particles must be properly oriented for bond breaking and bond making –for example, to be an effective collision between H 2 O and HCl, the oxygen of H 2 O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break

37. Energy Diagrams Energy diagram for an exothermic reaction

38. Energy Diagrams The reaction of H 2 and N 2 to form ammonia is exothermic –in this reaction, six covalent bonds are broken and six new ones are formed –breaking a bond requires energy, and forming a bond releases energy –in this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds; the reaction is exothermic

39. Energy Diagrams Energy diagram for an endothermic reaction

40. Factors Affecting Rate Nature of reactants –in general, reaction between ions in aqueous solution are very fast (activation energies are very low) –in general, reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher) Concentration –in most cases, reaction rate increases when the concentration of either or both reactants increases –for many reactions, there is a direct relationship between concentration and reaction rate; when concentration doubles the rate doubles

41. Factors Affecting Rate Temperature –in virtually all reactions, rate increases as temperature increases –an approximate rule for many reactions is that for a 10°C increase in temperature, the reaction rate doubles –when temperature increases, molecules move faster (have more kinetic energy), which means that they collide more frequently; more frequent collisions mean higher reaction rates –not only do molecules move faster at higher temperatures, but the fraction of molecules with energy equal to or greater than the activation energy also increases

42. Factors Affecting Rate The distribution of kinetic energies (molecular velocities) at two temperatures

43. Factors Affecting Rate Catalyst: Catalyst: a substance that increases the rate of a chemical reaction without itself being used up

44. Factors Affecting Rate Many catalysts provide a surface on which reactants can meet –the reaction of ethylene with hydrogen is an exothermic reaction –if these two reagents are mixed, there is no visible reaction even over long periods of time –when they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, the reaction takes place readily at room temperature

45. STOICIOMETRY The Quantities of Chemical Reactions

46. Formula Weight Formula weight: the sum of the atomic weights in atomic mass units (amu) of all atoms in a compound’s formula

47. Formula Weight formula weight formula weight can be used for both ionic and molecular compounds; it tells nothing about whether a compound is ionic or molecular molecular weight molecular weight should be used only for molecular compounds in this text, we use formula weight for ionic compounds and molecular weight for molecular compounds

48. The Mole Mole (mol) – a mole of the amount of substance that contains as many atoms, molecules, or ions as are in exactly 12 g of carbon-12 – a mole, whether it is a mole of iron atoms, a mole of methane molecules, or a mole of sodium ions, always contains the same number of formula units – the number of formula units in a mole is known as Avogadro’s number – Avogadro’s number has been measured experimentally – its value is 6.02214199 x 10 23 formula units per mole

49. Molar Mass Molar mass: Molar mass: the formula weight of a substance expressed in grams Glucose, C 6 H 12 O 6 – molecular weight: 180 amu – molar mass: 180 g/mol – one mole of glucose has a mass of 180 g Urea, (NH 2 ) 2 CO – molecular weight 60.0 amu – molar mass: 60.0 g/mol – one mole of urea has a mass of 60.0 g

50. Molar Mass We can use molar mass to convert from grams to moles, and from moles to grams – calculate the number of moles of water in 36.0 g water

51. Grams to Moles Calculate the number of moles of sodium ions, Na +, in 5.63 g of sodium sulfate, Na 2 SO 4 – first we find the how many moles of sodium sulfate – the formula weight of Na 2 SO 4 is 2(23.0) + 32.1 + 4(16.0) = 142.1 amu – therefore, 1 mol of Na 2 SO 4 = 142.1 g Na 2 SO 4 – the formula Na 2 SO 4 tells us there are two moles of Na + ions per mole of Na 2 SO 4

52. Grams to Molecules A tablet of aspirin, C 9 H 8 O 4, contains 0.360 g of aspirin. How many aspirin molecules is this? – first we find how many mol of aspirin are in 0.360 g – each mole of aspirin contains 6.02 x 10 23 molecules – the number of molecules of aspirin in the tablet is

53. Stoichiometry Stoichiometry: Stoichiometry: the study of mass relationships in chemical reactions – following is an overview of the the types of calculations we study

54. Stoichiometry Problem: how many grams of nitrogen, N 2, are required to produce 7.50 g of ammonia, NH 3 – first find how many moles of NH 3 are in 7.50 g of NH 3 – next find how many moles of N 2 are required to produce this many moles of NH 3

55. Stoichiometry Practice problem (cont’d) – finally convert moles of N 2 to grams of N 2 and now do the math

56. Stoichiometry Practice problems: – what mass of aluminum oxide is required to prepare 27 g of aluminum? 23 – how many grams each of CO 2 and NH 3 are produced from 0.83 mol of urea?

57. Limiting Reagent Limiting reagent Limiting reagent: the reagent that is used up first in a chemical reaction – consider this reaction of N 2 and O 2 – in this experiment, there is only enough O 2 to react with 1.0 mole of N 2 – O 2 is used up first; it the limiting reagent – 4.0 moles of N 2 remain unreacted

58. Limiting Reagent Practice Problem – suppose 12 g of carbon is mixed with 64 g of oxygen and the following reaction takes place – complete the following table. Which is the limiting reagent?

59. Percent Yield Actual yield: Actual yield: the mass of product formed in a chemical reaction Theoretical yield: Theoretical yield: the mass of product that should be formed according to the stoichiometry of the balanced chemical equation Percent yield: Percent yield: actual yield divided by theoretical yield times 100

60. Percent Yield Practice problem: – suppose we react 32.0 g of methanol with excess carbon monoxide and get 58.7 g of acetic acid – complete this table

61. End Chapter 5 Chemical Reactions