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Chapter 4 Chapter 4 Chemical Reactions Chemical reactions describe processes involving chemical change The chemical change involves rearranging matter Converting one or more pure substances into new pure substances Reactants Substances combined in the reaction Products Chemical Equations Chemical equations are used to describe chemical reactions The chemical symbols for the reactants are shown on the left The chemical symbols for the products are shown on the right An arrow (→) is used to indicate that reactants are converting to products A plus sign (+) is used to separate individual reactants and products Substances produced in the reaction Chapter 4 Chapter 4 Chemical Equations Chemical Equations “hydrogen and oxygen react to form water” Chemical equation describing the action: 2 H2(g) + O2(g) → 2 H2O(l) Subscripts? 2 H2(g) + O2(g) → 2 H2O(l) Or, P4(s) + 6Cl2 → 4PCl3 The equation accomplishes several things: 1. 2. Coefficients? Letters in parentheses? 3. Demonstrates the fundamental law of conservation of matter Atoms are neither created nor destroyed in chemical reactions, they are only rearranged Chapter 4 Balancing Chemical Equations Stoichiometry Pronounced as stoy-key-AHM-uh-tree For example, in the equation describing the formation of liquid water from hydrogen gas and oxygen gas 2 H2(g) + O2(g) → 2 H2O(l) There are four hydrogen atoms on both the left and right sides of the equation There are two oxygen atoms on both the left and right sides of the equation Therefore the equation is balanced “The relationship between the quantities of reactants and products” Chapter 4 Chapter 4 Balancing Chemical Equations Other examples NO(g) + O2(g) → NO2(g) Is it balanced? NO(g) + O(g) → NO2(g) Is this OK? How to Balance a Chemical Equation? NH3(g) + O2(g) → NO(g) +H2O (g) First step: balance nitrogen atoms on both sides NH3(g) + O2(g) → NO(g) +H2O (g) 2nd step: Then hydrogen atoms 2NH3(g) + O2(g) → NO(g) +3H2O (g) But I will need to balance the N atoms again 2NH3(g) + O2(g) → 2NO(g) +3H2O (g) 3rd step: Balance oxygens 2NH3(g) + (5/2) O2(g) → 2NO(g) +3H2O (g) Is it balanced? NO(g) + ½ O2(g) → NO2(g) Is this OK? 4th and final step: Just remove the fractions 4NH3(g) + 5 O2(g) → 4NO(g) +6H2O (g) Now try NO(g) + O2(g) → NO2(g) Chapter 4 Chapter 4 Balancing Chemical Equations H2S(aq) + I2(aq) → HI(aq) + S(s) Fe(s) + O2 → Fe2O3 KClO3(s) → KCl(s) + O2(g) Mg + O2(g) →MgO Ca(OH)2(s) + H3PO4(aq) → Ca3(PO4)2(s) +H2O(l) P4(s) +O2 (g) →P4O10(s) Ba(ClO3)2(aq) + H2SO4(aq) → HClO3(aq) + BaSO4(s) NH3(g)+O2(g) →NO(g) + H2O(g) C3H8(g) + O2(g) → CO2(g) + H2O(g) C8H18(l) + O2(g) → CO2(g) + H2O(g) Chapter 4 Balancing Chemical Equations H2S(aq) + I2(aq) → 2HI(aq) + S(s) 4Fe(s) + 3O2 → 2Fe2O3 2KClO3(s) → 2KCl(s) + 3O2(g) 2Mg + O2(g) →2MgO 3Ca(OH)2(s) + 2H3PO4(aq) → Ca3(PO4)2(s) +6H2O(l) P4(s) +5O2 (g) →P4O10(s) Ba(ClO3)2(aq) + H2SO4(aq) → 2HClO3(aq) + BaSO4(s) 4NH3(g)+5O2(g) →4NO(g) + 6H2O(g) C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) 2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g) Chapter 5 More Chemical Equations Balance the following equations SO2(g) + O2(g) → SO3(g) Types of Reactions Reaction Types Combination Decomposition Double Replacement Precipitation Acid-Base Single Replacement Oxidation-Reduction Gas Forming N2(g) + O2(g) → N2O(g) C6H14(l) + O2(g) → CO2(g) + H2O(l) HOMEWORK Chapter 5 Chapter 5 Combination Reactions Combination reactions have the form A+B→C Combination Reactions Other combination reactions: 2 Na(s) + S(s) → Na2S(s) Two or more reactants produce a single product SO3(g) + H2O(l) → H2SO4(aq) 2 H 2 + O2 =H 2 H2O 2Na (s) + Cl2 (g) → 2NaCl (s) =O Chapter 5 Chapter 5 Decomposition Reactions Decomposition reactions have the form A→B+C Single reactant breaks down into two or more products 2 H2O2 2 H2O + O2 =H =O Decomposition Reactions Further examples: 2 HgO(s) → 2 Hg(l) + O2(g) This reaction was used by Joseph Priestley in the discovery of oxygen in 1774 CaCO3(s) → CaO(s) + CO2(g) This reaction is used industrially to produce both lime (CaO) and CO2 from limestone (CaCO3) Air bags: 2 NaN3 (s) 2 Na (s) + 3 N2 (g) Chapter 5 Chapter 5 Double Replacement Double Replacement Double replacement reactions are also called “metathesis” reactions or “partner swapping” reactions They have the form AX + BY → BX + AY Double replacement reactions often take place in water and are of two basic types: Acid base neutralization reactions: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) (acid) (base) (salt) (water) H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l) + HF + =H NaOH =O → NaF = Na + H2O Precipitation reactions (solid forms): Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) + 2 NaNO3(aq) Here the barium sulfate precipitates out of solution =F Chapter 5 Chapter 5 Mixtures and Solutions Solute the component of a solution that is dissolved in another substance Mixtures and Solutions Electrolyte All ionic compounds that are soluble in water and conduct electricity Solvent the medium in which a solute dissolved to form a solution Strong electrolyte Solution a homogenous mixture in which the components are evenly distributed in each other Weak electrolyte Aqueous Any solution in which water is the solvent Precipitate Insoluble product of a reactant Non-electrolyte Chapter 5 Chapter 5 Solubility Rules Mixtures and Solutions Compounds Containing + + Group IA (Na , K ) and NH4 Nitrates Soluble Soluble Perchlorates (ClO4−) Soluble Chlorides, Bromides, Iodides Soluble CH3COOH(aq) CH3COO- (aq)+ H+ (aq) <5% ionized (weak electrolyte) Chapter 5 Exceptions Soluble Acetates (CH3COO−) CuCl2 (s) → Cu2+ (aq) + 2Cl- (aq) 100 % dissociation (strong electrolyte) Figure 5.2 Solubility + Ag+, Pb2+, Hg+ 2+ Fluorides Soluble Mg , Ca2+, Ag+, Pb2+, Sr2+ Sulfates Soluble Ag+, Ba2+, Sr2+, Pb2+, Hg+ Sulfides Insoluble Group IA and NH4+ Carbonates Insoluble Group IA and NH4+ Phosphates Insoluble Group IA and NH4+ Hydroxides Insoluble Group IA and NH4+ Chapter 5 Precipitation Reactions Precipitation reactions K2SO4 (aq) + Pb(NO3)2 (aq) → Ionic Equations When an ion appears on both sides of a chemical equation it can be canceled out “Spectator Ions” Writing Equations Molecular Net Ionic Equation Total ionic Net ionic http://www.dlt.ncssm.edu/TIGER/Flash/moles/DoubleDisp_Reaction-Precipitation.html Chapter 5 Chapter 5 Ionic Equations Ionic Equations Mg(ClO4)2(aq) + K2CO3(aq) → MgCO3(s) + KClO4(aq) Total: Cancel: KNO3(aq) + CaCl2(aq) → Total: Net Ionic: Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) + NaNO3(aq) Cancel: Net Ionic: Total: Cancel: Net ionic: Chapter 5 Chapter 5 Acids and Bases Acid a substance that, when dissolved in water, increases the concentration of hydrogen ions (H+) in solution Base a substance that, when dissolved in water, increases the concentration of hydroxide ions (OH-) in solution Strong acid Strong base Strong electrolyte Acids and Bases Chapter 5 Chapter 5 Acid-Base Reactions Acid-Base Reactions HNO3(aq) + NaOH(aq) → H2SO4(aq) + KOH(aq) → Total: Total: Cancel: Cancel: Net Ionic: Net Ionic: Chapter 5 Chapter 5 Single Replacement Single replacement reactions are also called substitution reactions They have the form A + BX → B + AX, where A and B are elements and BX and AX are compounds Single Replacement Other single replacement reactions 3 C(s) + 2 Fe2O3(s) → 4 Fe(s) + 3 CO2(g) Cu(NO3)2(aq) + Zn(s) → Zn(NO3)2(aq) + Cu(s) Single replacement reactions are also oxidationreduction (REDOX) reactions. Oxidation + Reduction H2 + CuO =H → =O H2O + Cu = Cu OXIDATION NUMBER Chapter 5 Chapter 5 Redox Reactions and Electron Transfer The oxidation number of an atom is the charge that atom would have if the compound was composed of ions. Oxidation of Mg (0) to Mg (II): 2Mg (s) + O2 (g) → 2MgO (s) Reduction of Fe (III) to Fe(0): Fe2O3 (s) + 3CO (g) →2Fe (s) + 3 CO2 (g) Reduction of Ag (I) to Ag and oxidation of Cu (0) to Cu (II): 2Ag+ (aq) + Cu(s) → 2 Ag (s) + Cu2+ (aq) Chapter 5 Gas Forming Reactions Gas Forming Reactions Reactions leading to the formation of an insoluble gas Examples Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + CO2 (g) + H2O (l)