1 Chapter 6 Chemical Reactions: An Introduction

2 Signs of a Chemical Reaction l Evolution of heat and light l Formation of a gas l Formation of a precipitate l Color change

3 Law of Conservation of Mass l mass is neither created nor destroyed in a chemical reaction 4 H 2 O 4 H 2 O 4 g32 g 36 g l total mass stays the same l atoms can only rearrange

4 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

5 In a chemical reaction l The way atoms are joined is changed l Atoms aren’t created of destroyed. l Can be described several ways l In a sentence l Copper reacts with chlorine to form copper (II) chloride. l In a word equation Copper + chlorine  copper (II) chloride

6 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

7 Symbols used in equations l (aq) after the formula - dissolved in water, an aqueous solution.  used after a product indicates a gas (same as (g))  used after a product indicates a solid (same as (s))

8 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.

9 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.

10 Diatomic elements l There are 8 elements that never want to be alone. l They form diatomic molecules. l H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, and At 2 l The –ogens and the –ines l pattern on the periodic table

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12 Skeleton Equation l Uses formulas and symbols to describe a reaction l doesn’t indicate how many. l All chemical equations are sentences that describe reactions.

13 Convert these to equations l Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (II) chloride and hydrogen sulfide gas. l Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

14 The other way Fe(g) + O 2 (g)  Fe 2 O 3 (s) Cu(s) + AgNO 3 (aq)  Ag(s) + Cu(NO 3 ) 2 (aq)

15 Balancing Chemical Equations

16 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.

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

18 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

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

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

21 Describing Equations l Describing Coefficients: –individual atom = “atom” –covalent substance = “molecule” –ionic substance = “unit” 3 molecules of carbon dioxide 2 atoms of magnesium 4 units of magnesium oxide 3CO 2  2Mg  4MgO 

22 Rules for balancing 1.Write the unbalanced equation. 2.Count atoms on each side. 3.Add coefficients to make #s equal. Coefficient  subscript = # of atoms 4.Reduce coefficients to lowest possible ratio, if necessary. 5.Double check atom balance!!!

23 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.

24 Helpful Tips l Balance one element at a time. l If an element appears more than once per side, balance it last. l Balance polyatomic ions as single units. –“1 SO 4 ” instead of “1 S” and “4 O” l If you fix everything except one element, and it is even on one side and odd on the other, double the first number, then move on from there. l C 4 H 10 + O 2  CO 2 + H 2 O

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

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

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

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

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

30 Example H 2 +H2OH2OO2O2  Recount RP H O

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

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

33 Examples CH 4 + O 2  CO 2 + H 2 O 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

34 Homework 14. a) Pb(NO 3 ) 2 + K 2 CrO 4  PbCrO 4 + KNO 3 b) MnO 2 + HCl  MnCl 2 + H 2 O+ Cl 2 c) C 3 H 6 + O 2  CO 2 +H 2 O d) Zn(OH) 2 + H 3 PO 4  Zn 3 (PO 4 ) 2 e) CO + Fe 2 O 3  Fe + CO 2 f) CS 2 + Cl 2  CCl 4 +S 2 Cl 2 g) CH 4 + Br 2  CH 3 Br + HBr h) Ba(CN) 2 + H 2 SO 4  BaSO 4 + HCN

35 Chapter 6 Summary

36 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.