1 Chemical Reactions

2 All chemical reactions l Have two parts l Reactants - the substances you start with l Products- the substances you end up with l The reactants turn into the products. Reactants  Products

3 In a chemical reaction l Law of Conservation of mass states: Atoms can not be created or destroyed. l Copper reacts with chlorine to form copper (II) chloride. l In a word equation: Copper + chlorine  copper (II) chloride l Cu + 2Cl → CuCl 2

4 Symbols used in equations l The arrow ( → ) separates the reactants from the products l Read “reacts to form” l The plus sign = “and” l (s) after the formula –solid l (g) after the formula -gas l (l) after the formula –liquid l (aq) after the formula - dissolved in water, an aqueous solution.

5 Symbols used in equations l indicates a reversible reaction (More later) l shows that heat is supplied to the reaction l is used to indicate a catalyst used supplied, in this case, platinum.

6 What is a catalyst? l A substance that speeds up a reaction without being changed by the reaction. l Enzymes are biological or protein catalysts.

7 Rates of Reactions In order for a reaction to occur: l The particles must collide (touch) l They must collide with enough energy l They must collide in the right orientation

8 Factors that affect reaction rate: l Temperature (particle energy) l Particle size l Surface area l Particle contact (stirring)

9 Balancing Chemical Equations

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

11 C + O 2  CO 2 l This equation is already balanced l What if it isn’t already? C + O O  C O O

12 C + O 2  CO l We need one more oxygen in the products. l Can’t change the formula, because it describes what is C + O  C O O

13 l Must be used to make another CO l But where did the other C come from? C + O  C O O O C

14 l Must have started with two C 2 C + O 2  2 CO C + O  C O O O C C

15 Rules for balancing  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)  Check to make sure it is balanced.

16 Never l Change a subscript to balance an equation. l If you change the formula you are describing a different reaction. l H 2 O is a different compound than H 2 O 2 l Never put a coefficient in the middle of a formula l 2 NaCl is okay, Na2Cl is not.

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

18 Example H 2 +H2OH2OO2O2  Need twice as much O in the product RP H O H O

19 Example H 2 +H2OH2OO2O2  Changes the O RP H O O H

20 Example H 2 +H2OH2OO2O2  Also changes the H RP H O H O

21 Example H 2 +H2OH2OO2O2  Need twice as much H in the reactant RP H O O H

22 Example H 2 +H2OH2OO2O2  Recount RP H O H O

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

24 Example H 2 +H2OH2OO2O2  This is the answer RP H O Not this

25 MINOH l A standard approach can be helpful in balancing chemical equations l The MINOH method provides the order in which elements will be considered when balancing an chemical equation –M – metals – I – polyatomic ions –N – non-metals (except oxygen and hydrogen) –O – oxygen –H - hydrogen

26 Balancing l __H 2 + __O 2  __H 2 O MINOHMINOH O - H - O - H

27 Balancing l __H 2 SO 4 + __NaOH  __H 2 O +__Na 2 SO 4 MINOHMINOH Na - SO 4 – O – H - Na - SO 4 – O – H

28 Balancing l __C 3 H 8 + __O 2  __CO 2 +__H 2 O MINOHMINOH C - O – H – C - O – H –

29 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

30 Evidence of Reactions Looking for the clues

31 Evidence of Reactions Just because the evidence is there DOES NOT mean a chemical reaction is taking place Have to look at everything that is going on

32 Evidence of Reactions 1.Formation of a gas Observations: Bubbles Smoke Odor/fumes

33 Evidence of Reactions 2.Change in color Be Careful! Some changes in color are physical changes!

34 Evidence of Reactions 3.Formation of a solid (precipitate) Observations: “Cloudy” “Foggy” Solid at bottom

35 Evidence of Reactions 4.Change in heat or light energy Observations: gets warm/hot/cold spark or explosion glows

36 Exothermic Reactions: Give off energy as heat or light WHY?? Because energy stored in chemical BONDS In EXOthermic reactions there is MORE energy stored in bonds of reactants than needed to form products! So.. there is left over energy!

37 Endothermic Reactions: Absorb (require) energy as heat or light WHY?? Because the energy stored in the bonds of the reactants is NOT ENOUGH to hold together the products. MORE ENERGY IS NEEDED!

38 Evidence of Reactions Summary 1.Formation of a gas 2.Change in color 3.Formation of a solid (precipitate) 4. Change in heat or light energy

39 Types of Reactions Predicting the Products

40 Types of Reactions l There are millions of reactions. l Can’t remember them all l Fall into several categories. l We will learn 5 types. l Will be able to predict the products. l For some we will be able to predict whether they will happen at all. l Will recognize them by the reactants

41 #1 Synthesis Reactions l Combine - put together l 2 elements, or compounds combine to make one compound. 2Ca +O 2  CaO SO 3 + H 2 O  H 2 SO 4

42 Write and balance Ca + Cl 2  Fe + O 2  iron (II) oxide Al + O 2  l Remember that the first step is to write the formula l Then balance

43 #2 Decomposition Reactions l decompose = fall apart l one reactant falls apart into two or more elements or compounds. l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2

44 #2 Decomposition Reactions l Can predict the products l Falls apart into its elements or small compounds lH2OlH2O l HgO

45 #2 Decomposition Reactions l NiCO 3 H 2 CO 3 (aq) 

46 #3 Single Replacement l One element replaces another l Reactants must be an element and a compound. l Products will be a different element and a different compound. Na + KCl  K + NaCl F 2 + LiCl  LiF + Cl 2

47 #4 Double Replacement l Two things replace each other. l Reactants must be two ionic compounds or acids. NaOH + FeCl 3  l The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 NaOH + FeCl 3  Fe(OH) 3 + NaCl

48 How to recognize which type l Look at the reactants l One Product = Synthesis l One Reactant = Decomposition l E + CSingle Replacement l C + CDouble Replacement

49 Examples H 2 + O 2  H 2 O  Zn + H 2 SO 4  HgO  KBr +Cl 2  AgNO 3 + NaCl  Mg(OH) 2 + H 2 SO 3 

50 Last Type l Combustion of hydrocarbons l A compound composed of only C H and maybe O is reacted with oxygen l If the combustion is complete, the products will be CO 2 and H 2 O.

51 Examples C 4 H 10 + O 2  CO 2 and H 2 O C 6 H 12 O 6 + O 2  CO 2 and H 2 O

52 Chapter 6 Summary

53 An equation l Describes a reaction l Must be balanced because to follow Law of Conservation of Energy l Can only be balanced by changing the coefficients. l Has special symbols to indicate state, and if catalyst or energy is required.

54 Reactions l Come in 5 types. l Can tell what type they are by the reactants.

55 The Process l Determine the type by looking at the reactants. l Put the pieces next to each other l Use charges to write the formulas l Use coefficients to balance the equation.