Chemical Equations Chapter 8 Hein and Arena Version 2.0 12th Edition Eugene Passer Chemistry Department Bronx Community 1 College © John Wiley and Sons, Inc Chapter Outline 8.1 The Chemical Equation 8.4 Types of Chemical Equations 8.2 Writing and Balancing Chemical Equations 8.5 Heat in Chemical Reactions 8.3 Information in a Chemical Equation 2 • Chemists use chemical equations to describe reactions they observe in the laboratory or in nature. • Chemical equations provide us with the means to 1. summarize the reaction 2. display the substances that are reacting 3. show the products 4. indicate the amounts of all component substances in a reaction. 3 8.1 The Chemical Equation 4 • Chemical reactions always involve change. • Atoms, molecules or ions rearrange to form different substances. • The substances entering the reaction are called reactants. • The substances formed in the reaction are called products. • During reactions, chemical bonds are broken and new bonds are formed. 5 • A chemical equation is a shorthand expression for a chemical change or reaction. • A chemical equation uses the chemical symbols and formulas of the reactants and products and other symbolic terms to represent a chemical reaction. 6 Chemical Equation Al + Fe2O3 Fe + Al2O3 reactants products iron oxygen bonds break aluminum oxygen bonds form 7 Coefficients (whole numbers) are placed in front of substances to balance the equation and to indicate the number of units (atoms, molecules, moles, or ions) of each substance that are reacting. 8 2 Al + Fe2O3 2 Fe + Al2O3 coefficient coefficient 9 Conditions required to carry out the reaction may be placed above or below the arrow. 10 2 Al + Fe2O3 2 Fe + Al2O3 coefficient coefficient heat 11 The physical state of a substance is indicated by symbols such as (l) for liquid. 12 In a chemical reaction atoms are neither created nor destroyed. 2Al(s) (s) + Fe2O3(s) 2Fe(l) (s) + Al2O3 (s) All atoms present in the reactant must also be present in the products. 13 Symbols Used in Chemical Reactions 14 symbol + meaning plus location placed between substances 15 symbol meaning yields location between reactants and products 16 symbol (s) meaning solid location after formula 17 symbol (l) meaning liquid location after formula 18 symbol (g) meaning gas location after formula 19 symbol (aq) meaning aqueous location after formula 20 symbol meaning heat location written above or below 21 symbol h meaning light energy location written above 22 symbol meaning gas formation location after formula 23 8.2 Writing and Balancing Equations 24 • To balance an equation adjust the number of atoms of each element so that they are the same on each side of the equation. • Never change a correct formula to balance an equation. 25 Steps for Balancing Equations 26 Step 1 Identify the reaction. Write a description or word equation for the reaction. Mercury (II) oxide decomposes to form mercury and oxygen. mercury(II) oxide → mercury + oxygen 27 Step 2 Write the unbalanced (skeleton) equation. – The formulas of the reactants and products must be correct. – The reactants are written to the left of the arrow and the products to the right of the arrow. HgO Hg + O2 The formulas of the reactants and products can never be changed. 28 Step 3a Balance the equation. – Count and compare the number of atoms of each element on both sides of the equation. – Determine the elements that require balancing. 29 Step 3a Balance the equation. HgO → Hg + O2 Element Hg Reactant Side 1 Product Side 1 – There is one mercury atom on the reactant side and one mercury atom on the product side. – Mercury is balanced. 30 Step 3a Balance the equation. HgO Hg + O2 Element O Reactant Side 1 Product Side 2 – There are two oxygen atoms on the product side and there is one oxygen atom on the reactant side. – Oxygen needs to be balanced. 31 Step 3b Balance the equation. – Balance each element one at a time, by placing whole numbers (coefficients) in front of the formulas containing the unbalanced element. – A coefficient placed before a formula multiplies every atom in the formula by that coefficient. 32 Step 3b Balance the equation. 2HgO Hg + O2 Element O Reactant Side 12 Product Side 2 – Place a 2 in front of HgO to balance O. – There are two oxygen atoms on the reactant side and there are two oxygen atoms on the product side. – Oxygen (O) is balanced. 33 Step 3c Balance the equation. – Check all other elements after each individual element is balanced to see whether, in balancing one element, another element became unbalanced. 34 Step 3c Balance the equation. 2HgO Hg + O2 Element Hg Reactant Side 2 Product Side 1 – Count and compare the number of mercury (Hg) atoms on both sides of the equation. – There are two mercury atoms on the reactant side and there is one mercury atom on the product side. – Mercury (Hg) is not balanced. 35 Step 3c Balance the equation. 2HgO 2Hg + O2 Element Hg Reactant Side 2 Product Side 12 – Place a 2 in front of Hg to balance mercury. – There are two mercury atoms on the reactant side and there are two mercury atoms on the product side. – Mercury (Hg) is balanced. 36 THE EQUATION IS BALANCED 2HgO 2Hg + O2 Element Hg O Reactant Side 2 2 Product Side 2 2 37 Balance the Equation sulfuric acid + sodium hydroxide → sodium sulfate + water 38 Balance the Equation H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + H2O(l) Reactant Side 1 Product Side 1 Na 12 2 O 12 1 H 34 2 SO 24 Place ais2one There in front Na on ofthe NaOH reactant to balance side and Na.there are two Na on the product side. 39 THE EQUATION IS BALANCED 2 + H2SO4(aq)2 + NaOH(aq) → Na2SO4(aq) H2O(l) Reactant Side Product Side 2SO 4 1 1 Na 2 2 O 2 12 H 4 42 There 4H on the side and two H on Place aare 2 in front of Hreactant 2O to balance H. the product side. 40 Balance the Equation butane + oxygen → carbon dioxide + water 41 Balance the Equation C4H10 (g) + O2 (g) → 4 CO2(g) + H2O(l) C Reactant Side 4 Product Side 14 H 10 2 O 2 39 Place aare There 4 in four front C on of the CO2reactant to balance sideC. and there is one C on the product side. 42 C4H10 (g) + O2 (g) → 4 CO2(g) + 5H2O(l) C Reactant Side 4 Product Side 4 H 10 2 10 O 2 9 13 Place aare There 5 in 10front H onofthe H2reactant O to balance side and H. there are two H on the product side. 43 10 2 C4H10 (g) + O2 (g) → 48 CO2(g) +10 55 H2O(l) C Reactant Side 48 Product Side 84 H 10 20 10 20 O 2 13 26 To balance There is no O, whole double number the coefficients coefficient of that each can be placed in front substance otherofthan O2 oxygen. to balance O. 44 THE EQUATION IS BALANCED 2C4H10 (g) +13O2 (g) → 8CO2(g) +10H2O(l) C Reactant Side 8 Product Side 8 H 20 20 O 2 26 26 There aare Place 13now in front 26 Oofon O2the to product balance side. O. 45 8.3 Information in a Chemical Equation 46 The meaning of a formula is context dependent. The formula H2O can mean: 1. 2. 3. 4. 5. 2H and 1 O atom 1 molecule of water 1 mol of water 6.022 x 1023 molecules of water 18.02 g of water 47 In an equation, formulas can represent units of individual chemical entities or moles. H2 + 1 molecule H2 1 mol H2 Cl2 1 molecule Cl2 1 mol Cl2 → 2HCl 2 molecules HCl 2 mol HCl 48 Formulas Number of molecules Number of atoms Number of moles Molar masses 49 8.4 Types of Chemical Equations 50 Combination Decomposition Single-Displacement Double-Displacement 51 Combination Reactions 52 Two reactants combine to form one product. A + B AB 53 Examples 54 Metal + Oxygen → Metal Oxide 2Ca(s) + O2(g) 2CaO(s) 4Al(s) + 3O2(g) 2Al2O3(s) 55 Nonmetal + Oxygen → Nonmetal Oxide S(s) + O2(g) SO2(g) N2(g) + O2(g) 2NO(g) 56 Metal + Nonmetal → Salt 2K(s) + F2(g) 2KF(s) 2Al(s) + 3Cl2(g) 2AlCl3(s) 57 Metal Oxide + Water → Metal Hydroxide Na2O(s) + H2O(l) 2NaOH(aq) CaO(s) + 2H2O(l) 2Ca(OH)2(aq) 58 Nonmetal Oxide + H2O(l) → Oxy-acid SO3(g) + H2O(l) H2SO4(aq) N2O5(g) + H2O(l) 2HNO3(aq) 59 Decomposition Reactions 60 A single substance breaks down to give two or more different substances. AB A + B 61 Examples 62 Metal Oxide → Metal + Oxygen 2Ag2O(s) 4Ag(s) + O2(g) Metal Oxide → Metal Oxide + Oxygen 2PbO2(s) 2PbO(s) + O2(g) 63 Carbonate → CO2(g) CaCO3(s) CaO(s) + CO2(g) Hydrogen Carbonate → CO2(g) 2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g) 64 Miscellaneous Reactions 2H2O2(l) 2H2O(l) + O2(g) 2KClO3(s) 2KCl(s) + 3O2(g) 2NaNO3(s) 2NaNO2(s) + O2(g) 65 Single Displacement Reactions 66 One element reacts with a compound to replace one of the elements of that compound. A + BC AC + B 67 Metal + Acid → Hydrogen + Salt Mg(s) + 2HCl(aq) H2(g) + MgCl2(aq) salt 2Al(s) + 3H2SO4(aq) 3H2(g) + Al2(SO4)3(aq) salt 68 Metal + Water → Hydrogen + Metal Hydroxide 2Na(s) + 2H2O(l) H2(g) + 2NaOH(aq) metal hydroxide Ca(s) + 2H2O(l) H2(g) + Ca(OH)2(aq) metal hydroxide 69 Metal + Water → Hydrogen + Metal Oxide 3Fe(s) + 4H2O(g) 4H2(g) + Fe3O4(s) metal oxide 70 The Activity Series 71 An atom of an element in the activity series will displace an atom of an element below it from one of its compounds . increasing activity Metals K Ca Na Mg Al Zn Sodium (Na) will displace an Fe atom below it from one of its Ni Sn compounds. Pb H Cu Ag Hg 72 Examples Metal Activity Series 73 Metal Higher in Activity Series Displacing Metal Below It Mg(s) + PbS(s) MgS(s) + Pb(s) Metals Mg Al Zn Fe Ni Sn Pb Magnesium is above lead in the activity series. 74 Metal Lower in Activity Cannot Displace Metal Above It Ag(s) + CuCl2(s) no reaction Metals Pb H Cu Ag Hg Silver is below copper in the activity series. 75 Example Halogen Activity Series 76 Halogen Higher in Activity Series Displaces Halogen Below It Cl2(g) + CaBr2(s) CaCl2(aq) + Br2(aq) Halogens F2 Cl2 Br2 I2 Chlorine is above bromine in the activity series. 77 Double Displacement Reactions 78 Two compounds exchange each The reaction can be thoughtpartners of as anwith exchange A B displaces C D and combines with D C other to produce two different compounds. of positive and negative groups. AB + CD AD + CB 79 The Following Accompany Double Displacement Reactions • • • • formation of a precipitate release of gas bubbles release of heat formation of water 80 Examples 81 Acid Base Neutralization acid + base → salt + water HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l) 82 Formation of an Insoluble Precipitate AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq) 83 Metal Oxide + Acid metal oxide + acid → salt + water CuO(s) + 2HNO3(aq) Cu(NO3)2(aq) + H2O(l) CaO(s) + 2HCl(aq) CaCl2(s) + H2O(l) 84 Formation of a Gas H2SO4(aq) + 2NaCN(aq) Na2SO4(aq) + 2HCN(g) NH4Cl(aq) + NaOH(aq) NaCl(aq) + NH4OH(aq) indirect gas formation NH4OH(aq) NH3(g) + H2O(l) 85 8.5 Heat in Chemical Reactions 86 Energy changes always accompany chemical reactions. Whenreason One this why occurs, reactions energy occur is is that thetoproduct released the surroundings. attains a lower energy state than the reactants. 87 The amounts of substances areliberate expressed in moles. Exothermic reactions heat. H2(g) + Cl2(g) → 2HCl(g) + 185 kJ (exothermic) 1 mol 1 mol 2 mol Endothermic reactions absorb heat. N2(g) + O2(g) + 185 kJ → 2NO(g) (exothermic) 1 mol 1 mol 2 mol 88 For life on Earth the sun is the major provider of energy. The energy for plant photosynthesis is derived from the sun. 6CO2 + 6H2O + 2519 kJ → C6H12O6 + 6O2 glucose glucose 89 Energy of Activation 90 • A certain amount of energy is always required for a reaction to occur. • The energy required to start a reaction is called the energy of activation. 91 CH4 + 2O2 → CO2 + 2H2O + 890 kJ • This reaction will not occur unless activation energy is supplied. • The activation energy can take the form of a spark or a flame. 92 93 8.2 8.3 94