1. Chemical Reactions Dr. Schuerch

2. Describing Chemical Reactions A chemical reaction occurs when one or more reactants changes into one or more products, example: H 2 O 2  H 2 O + O 2 H 2 O 2 is the reactant H 2 O and O 2 are the products + used to separate two reactants or two products  “Yields,” separates reactants from products

3. Symbols Used in Chemical Equations + (s) (l) (g) (aq) Used to separate two reactants or two products “Yields,” separates reactants from products Used in place of  for reversible reactions Designates a reactant or product in the solid state; placed after the formula Designates a reactant or product in the liquid state: placed after the formula Designates a reactant or product in the gaseous state; placed after the formula Designates an aqueous solution; the substance is dissolved in water; placed after the formula Indicates that heat is supplied to the reaction A formula written above or below the yield sign indicates its use as a catalyst (in this example, platinum). Heat Δ Pt

4. Word Equations Describe what happens in a chemical reaction –To write a word equation, write the names of the reactants to the left of the arrow separated by plus signs; write the names of the products to the right of the arrow, also separated by a plus sign –Examples: Reactants  Products Iron + oxygen  iron(III) oxide Hydrogen peroxide  water + oxygen Methane + oxygen  carbon dioxide + water

5. Chemical Equations Chemical equations are a representation of a chemical reaction: the formulas of the reactants on the left are connected by an arrow with the formulas for the products on the right –Example: Fe + O 2  Fe 2 O 3 Indicate the physical state of reactants and products and if a catalyst is used –Example: Fe (s) + O 2(g)  Fe 2 O 3(s) Balance the Equation –Example: 4Fe (s) + 3O 2(g)  2Fe 2 O 3(s)

6. Balancing Chemical Equations To write a balanced chemical equation, first write the skeleton equation. Then use the coefficients to balance the equation so that it obeys the law of conservation of mass –Example: H 2 O 2  O 2 + H 2 O (Skeleton Equation) 2H 2 O 2 (aq)  O 2 (g) + 2H 2 O(l) (Balanced Equation) Coefficients

7. Rules for Writing and Balancing Equations 1.Determine the correct formulas for all the reactants and products 2.Write the skeleton equation by placing the formulas for the reactants on the left and the formulas for the products on the right with a yield sign (  ) in between. If two or more reactants or products are involved, separate their formulas with plus signs 3.Determine the number of atoms of each element in the reactants and products. Count a polyatomic ion as a single unit if it appears unchanged on both sides of the equation 4.Balance the elements one at a time by using coefficients. When no coefficient is written, it is assumed to be 1. Begin by balancing elements that appear only once on each side of the equation. Never balance an equation by changing the subscripts in a chemical formula. 5.Check each atom or polyatomic ion to be sure they are equal on both sides of the equation

8. Balance the Following Equations 1. Al(s) + O 2 (g)  Al 2 O 3 (s) 2. AgNO 3 + H 2 S  Ag 2 S + HNO 3 3. Zn(OH) 2 + H 3 PO 4  Zn 3 (PO 4 ) 2 + H 2 O 4. FeCl 3 + NaOH  Fe(OH) 3 + NaCl

9. Predicting whether a reaction will occur Driving forces of a reaction –Transfer of electrons –Formation of a solid –Formation of water –Formation of a gas

10. Types of Chemical Reactions A.Oxidation - Reduction Reactions 1.Combination or synthesis Reactions 2.Decomposition Reactions 3.Single-Replacement Reactions 4.Combustion Reactions B.Double-Replacement Reactions 1.Precipitation Reactions 2.Acid-base Reactions

11. Oxidation-Reduction Combination aka Synthesis Reactions A chemical change in which two or more substances react to form a single new substance –Examples 2Na(s) + Cl 2 (g)  2NaCl(s) S(s) + O 2 (g)  SO 2 (g) 2S(s) + 3O 2 (g)  2SO 3 (g) Fe(s) + S(s)  FeS(s) 2Fe(s) + 3S(s)  Fe 2 S 3 (s) Note: When two non metals or a transition metal and non metal are present in the reaction, more than one product may result

12. Oxidation-Reduction Decomposition Reaction Chemical reactions in which a single compound breaks down into two or more simpler products –Examples: 2HgO(s)  2Hg(l) + O 2 (g) 2H 2 O(l)  2H 2 (g) + O 2 (g) 2HI  H 2 (g) + I 2 (s) Δ Δ Electricity

13. Oxidation-Reduction Single Replacement Reactions A chemical change in which one element replaces a second element in a compound –Identified by both reactants and products having an element and a compound –Examples: 2K(s) + 2HOH(L)  2KOH(aq) + H 2 (g) Br 2 (aq) + 2NaI(aq)  2NaBr + I 2 (aq) Zn(s) + Cu(NO 3 ) 2 (aq)  Cu(s) + Zn(NO 3 ) 2 (aq)

14. Single Replacement Reactions Whether one metal will displace another metal from a compound depends upon the relative reactivities of the two metals –Example Zn(s) + Cu(NO 3 ) 2 (aq)  Cu(s) + Zn(NO 3 ) 2 (aq) Cu(s) + Zn(NO 3 ) 2 (aq)  No reaction –Zinc can displace copper because it is more reactive than copper –Copper cannot displace zinc because it is less reactive than zinc

15. Activity Series of Metals Notice that zinc is higher on the activity series of metals than copper A reactive metal will replace any metal listed below it in the activity series Zn + Cu(NO 3 ) 2  Cu + Zn(NO 3 ) 2 Cu + Zn(NO 3 ) 2  No reaction Decreasing reactivity Oxidation Reduction Single Replacement Reactions

16. Halogens (Group 7) also participate in single replacement reactions –Example: Br 2 + 2NaI  2NaBr + I 2 NaBr + I 2  No reaction Halogens higher in the periodic table will replace halogens lower in the periodic table Oxidation Reduction Single Replacement Reactions Decreasing reactivity

17. Oxidation Reduction Combustion Reactions Is a chemical change in which an element or a compound reacts with oxygen often producing energy of the form of heat and light –Examples: 2C 8 H 16 (l) + 25O 2 (g)  16CO 2 (g) + 18H 2 O(l) 2Mg(s) + O 2 (g)  2MgO(s) S(s) + O 2 (g)  SO 2 (g) –Note that some of the combustion reactions are also combination reactions

18. Oxidation-Reduction Reaction

19. Double Replacement Reactions A chemical change involving an exchange of positive ions between two compounds –Generally occur in aqueous solution and often produce a precipitate, a gas, or molecular compound such as water –Examples: –Precipitation (forms a solid or gas as a product) Na 2 S(aq) +Cd(NO 3 ) 2 (aq)  CdS(s) + 2NaNO 3 (aq) 2NaCN(aq) + H 2 SO 4 (aq)  2HCN(g) + Na 2 SO 4 (aq) –Acid Base (forms water as product) Ca(OH) 2 (aq) + 2HCl(aq)  CaCl 2 (aq) + 2H 2 O(l)

20. Predicting the Formation of a Precipitate You can predict the formation of a precipitate by using the general rules for solubility of ionic compounds (Table 7.1 pg 178 in your text)

21. Reactions in Aqueous Solution: Complete Ionic Equations Many important reactions take place in water—that is in aqueous solution –When in solution, many ionic compounds (called strong electrolytes) will ionize splitting into their constitute ions Weak electrolytes are ionic compounds that are only slightly soluble in water and when formed in water precipitate out of solution (we can think of them as being insoluble) –Equations in aqueous solution can be written as complete ionic equations Example: A molecular equation AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq) Can be written as a complete ionic equation: Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) + NO 3 - (aq)

22. Reactions in Aqueous Solution: Net Ionic Equations Given the complete ionic equation for the formation of silver chloride precipitate from the mixture of silver nitrate and sodium chloride Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) + NO 3 - (aq) –Notice that the sodium and nitrate ion appear unchanged on both sides of the equation Because they are unchanged in the chemical reaction they are called spectator ions –When you rewrite the ionic equation without the spectator ions, you get a net ionic equation Ag + (aq) + Cl - (aq)  AgCl(s) A net ionic equation shows only those particles involved in the reaction and is balanced with respect to both mass and charge

23. Reactions in Aqueous Solution: Net Ionic Equations (Continued) Consider the following skeleton equation Pb(s) + AgNO 3 (aq)  Ag(s) + Pb(NO 3 ) 2 (aq) –Write a complete ionic equation Pb(s) + Ag + (aq) + NO 3 - (aq)  Ag(s) + Pb 2+ (aq) 2NO 3 - (aq) –Write a net ionic equation Pb(s) + Ag + (aq)  Ag(s) + Pb 2+ (aq) –Is the net ionic equation balanced (Notice the charge on both sides of the equation)? Balance the net ionic equation Pb(s) + 2Ag + (aq)  2Ag(s) + Pb 2+ (aq)

24. Reactions in Aqueous Solution Acid-Base reactions Acid - a substance that produces H+ ions when dissolved in water (Arrhenius definition) –Strong acids completely disassociate in solution and can be considered a strong electrolyte Base - is a substance that produces OH- ions when dissolved in solution (Arrhenius definition) –Strong bases completely disassociate in solution and can be considered a strong electrolyte

25. Reactions in Aqueous Solution Acid-Base reactions Acids and Bases react to form water and a ionic compound called a salt