Chemical Reactions Ch. 4

1. Chemical ReactionsCh. 4 Milbank High School

2. Sec. 4.1Balancing Chemical Equations • Objectives • Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction.

3. Balanced Equation • Atoms can’t be created or destroyed • All the atoms we start with we must end up with • A balanced equation has the same number of each element on both sides of the equation.

4. Rules for balancing: • Assemble, write the correct formulas for all the reactants and products • Count the number of atoms of each type appearing on both sides • Balance the elements one at a time by adding coefficients (the numbers in front) - save H and O until LAST! • Check to make sure it is balanced.

5. Never • Never change a subscript to balance an equation. • If you change the formula you are describing a different reaction. • H2O is a different compound than H2O2 • Never put a coefficient in the middle of a formula • 2 NaCl is okay, Na2Cl is not.

6. Example H2 + O2 ® H2O Make a table to keep track of where you are at

7. Example H2 + O2 ® H2O R P 2 H 2 2 O 1 Need twice as much O in the product

8. Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 Changes the O

9. Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 2 Also changes the H

10. Example H2 + O2 ® 2 H2O R P 2 H 2 4 2 O 1 2 Need twice as much H in the reactant

11. Example 2 H2 + O2 ® 2 H2O R P 2 H 2 4 2 O 1 2 Recount

12. Example 2 H2 + O2 ® 2 H2O R P 4 2 H 2 4 2 O 1 2 The equation is balanced, has the same number of each kind of atom on both sides

13. Example 2 H2 + O2 ® 2 H2O R P 4 2 H 2 4 2 O 1 2 This is the answer Not this

14. Balancing Examples • _AgNO3 + _Cu ® _Cu(NO3)2 + _Ag • _Mg + _N2® _Mg3N2 • _P + _O2® _P4O10 • _Na + _H2O ® _H2 + _NaOH • _CH4 + _O2® _CO2 + _H2O

15. Sec. 4.2 Volume Relationships in Chemical Equations • Objectives • Define Gay-Lussac’s Law • Determine how to find volume relationships in a gaseous chemical reaction

16. Gay-Lussac’s Law • When gases measured at the same temperature and pressure are allowed to react, the volumes of gaseous reactants and products are in small whole-number ratios • 3 hydrogen gas volumes + 1 nitrogen gas volumes = 2 volumes ammonia • (3:1:2) • Practice Exercises Pg. 103

17. Sec. 4.3Avogadro’s Number • Objectives • Define Avogadro’s Number

18. Moles (abbreviated: mol) • Defined as the number of carbon atoms in exactly 12 grams of carbon-12. • 1 mole is 6.02 x 1023 particles. • 6.02 x 1023 is called Avogadro’s number. • How large is it?

19. Sec. 4.4Molecular Masses and Formula Masses • Objectives • Define molecular and formula mass

20. Molecular Mass • Molecular mass is the average mass of a molecule of a substance relative to that of a carbon-12 atom • sum of masses of the atoms represented in a molecular formula • Formula mass is the same as the molar mass except that it refers to ionic compounds

21. Sec. 4.5Chemical Arithmetic and the Mole • Objectives • Define and determine molar mass

22. Molar Mass • Molar mass is the generic term for the mass of one mole of any substance (in grams) • Numerically equal to atomic mass, molecular mass, or formula mass • Expressed as g/mol

23. Examples • Calculate the molar mass of the following and tell what type it is: • Na2S • N2O4 • C • Ca(NO3)2 • C6H12O6 • (NH4)3PO4

24. What is a Mole? • You can measure mass, • or volume, • or you can count pieces. • We measure mass in grams. • We measure volume in liters. • We count pieces inMOLES.

25. Molar Mass • The number of grams of 1 mole of atoms, ions, or molecules. • We can make conversion factors from these. • To change grams of a compound to moles of a compound.

26. For example • How many moles is 5.69 g of NaOH?

27. For example • How many moles is 5.69 g of NaOH?

28. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles

29. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH

30. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g

31. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

32. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

33. For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g

34. Examples • How many moles is 4.56 g of CO2? • How many grams is 9.87 moles of H2O? • How many molecules is 6.8 g of CH4? • 49 molecules of C6H12O6 weighs how much?

35. 4.6 Mole and Mass Relationships in Chemical Equations • Objectives • Determine what a stoichiometric factor is • Solve problems using stoichiometric factors

36. Stoichiometric Factors • Relates the amounts of any two substances involved in chemical equations

37. Rules • Write a balanced equation • Determine molar masses • Use molar mass to convert grams to moles of substance • Convert moles of given substance to moles of the substance you wish to find • Use molar mass to convert moles to grams of that substance

38. Example 4.11 • Example 4.12 • Practice Exercises

39. Limiting Reagent • Limits or determines the amount of product that can be formed in a reaction • Excess reagent • Reactant that is not completely used up

40. Example • Sodium chloride can be prepared by the reaction of sodium metal with chlorine gas. 2Na + Cl2 2NaCl Suppose that 6.70 mol Na reacts with 3.20 mol Cl2. What is the limiting reagent? How many moles of NaCl are produced?

41. Sec. 4.7Structure, Stability, and Spontaneity • Objectives • Determine what exothermic and endothermic reactions are • Define enthalpy

42. Heat released or absorbed • Heat of reaction • Heat released or absorbed during a chemical reaction • Enthalpy change • Equal to heat of reaction • ∆H • Exothermic • Release heat, -∆H • Endothermic • Absorb heat, ∆H

43. Sec. 4.8Forward and Reverse Reactions • Objectives • Define activation energy • Recognize the components of a reaction profile

44. Activation Energy • Minimum energy needed to get the reaction started • AKA energy of activation • Ea • See reaction profiles

45. Sec. 4.9Reaction Rates: Collisions and Orientation • Reaction rates profoundly affected by temperature, catalysts, and concentration • For a reaction to occur: • Collision of particles • Proper orientation (Fig. 4.13) • Minimum activation energy

46. Temperature • Generally faster reactions at higher temps • More collisions, more energy • Uses: • Freezing food • Cooking • Organisms • Fevers • Hibernation • Heart surgery

47. Catalysts • A substance that changes the rate of reaction without being changed itself • Lowers the activation energy (Fig. 4.15) • Increase speed of slow reaction • Enzymes • Biological catalysts • Mediate reactions in living systems

48. Concentration • More molecules = more collisions • Double concentration = double rate of reaction

49. Mechanism • Step-by-step process of a reaction • Affected by temp, catalysts, and concentration • Cancer • Break-down in the mechanism of a reaction • Poisons • Knowledge of mechanisms has led to effective treatments

50. Sec. 4.10Equilibrium in Chemical Reactions • Some reactions proceed in forward and reverse directions at the same time • Double arrow • Equilibrium eventually reached • Rates of forward reaction = rates of reverse reaction