1. Chemical Reactions Ch. 4 Milbank High School

2. Sec. 4.1 Balancing Chemical Equations  Objectives Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction. 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 change a subscript to balance an equation. If you change the formula you are describing a different reaction. If you change the formula you are describing a different reaction. H 2 O is a different compound than H 2 O 2 H 2 O is a different compound than H 2 O 2  Never put a coefficient in the middle of a formula 2 NaCl is okay, Na2Cl is not. 2 NaCl is okay, Na2Cl is not.

6. Example H 2 +H2OH2OO2O2  Make a table to keep track of where you are at

7. Example H 2 +H2OH2OO2O2  Need twice as much O in the product RP H O 2 2 2 1

8. Example H 2 +H2OH2OO2O2  Changes the O RP H O 2 2 2 1 2

9. Example H 2 +H2OH2OO2O2  Also changes the H RP H O 2 2 2 1 2 2

10. Example H 2 +H2OH2OO2O2  Need twice as much H in the reactant RP H O 2 2 2 1 2 2 4

11. Example H 2 +H2OH2OO2O2  Recount RP H O 2 2 2 1 2 2 4 2

12. Example H 2 +H2OH2OO2O2  The equation is balanced, has the same number of each kind of atom on both sides RP H O 2 2 2 1 2 2 4 2 4

13. Example H 2 +H2OH2OO2O2  This is the answer RP H O 2 2 2 1 2 2 4 2 4 Not this

14. Balancing Examples  _ AgNO 3 + _Cu  _Cu(NO 3 ) 2 + _Ag  _Mg + _N 2  _Mg 3 N 2  _P + _O 2  _P 4 O 10  _Na + _H 2 O  _H 2 + _NaOH  _CH 4 + _O 2  _CO 2 + _H 2 O

15. Sec. 4.2 Volume Relationships in Chemical Equations  Objectives Define Gay-Lussac’s Law Define Gay-Lussac’s Law Determine how to find volume relationships in a gaseous chemical reaction 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. 4.3 Avogadro’s Number  Objectives Define Avogadro’s Number 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 10 23 particles.  6.02 x 10 23 is called Avogadro’s number.  How large is it?

19. Sec. 4.4 Molecular Masses and Formula Masses  Objectives Define molecular and formula mass 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 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. 4.5 Chemical Arithmetic and the Mole  Objectives Define and determine molar mass 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 Numerically equal to atomic mass, molecular mass, or formula mass Expressed as g/mol Expressed as g/mol

23. Examples  Calculate the molar mass of the following and tell what type it is:  Na 2 S N2O4N2O4N2O4N2O4 CCCC  Ca(NO 3 ) 2  C 6 H 12 O 6  (NH 4 ) 3 PO 4

24. What is a Mole? 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 in MOLES.

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. 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? l need to change grams to moles

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

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

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

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

34. Examples  How many moles is 4.56 g of CO 2 ?  How many grams is 9.87 moles of H 2 O?  How many molecules is 6.8 g of CH 4 ?  49 molecules of C 6 H 12 O 6 weighs how much?

35. 4.6 Mole and Mass Relationships in Chemical Equations  Objectives Determine what a stoichiometric factor is Determine what a stoichiometric factor is Solve problems using stoichiometric factors 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 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 Reactant that is not completely used up

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

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

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

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

44. Activation Energy  Minimum energy needed to get the reaction started AKA energy of activation AKA energy of activation E a E a See reaction profiles See reaction profiles

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

46. Temperature  Generally faster reactions at higher temps More collisions, more energy More collisions, more energy  Uses: Freezing food Freezing food Cooking Cooking Organisms Organisms FeversFevers HibernationHibernation Heart surgeryHeart 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 Biological catalysts Mediate reactions in living systems 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 Break-down in the mechanism of a reaction  Poisons Knowledge of mechanisms has led to effective treatments Knowledge of mechanisms has led to effective treatments

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

51. La Chatelier’s Principle  If a stress if applied to a system, it rearranges itself to relieve the stress  Changes in concentration, temp, pressure  Catalysts do not change the equilibrium