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Chapter 6 Chemical Reactions Chemical Reactions In a chemical reaction, one or more reactants is converted to one or more products Reactant(s) Product(s) In this chapter we discuss three aspects of chemical reactions (a) mass relationships (stoichiometry) (b) types of reactions (c) heat gain and loss accompanying reactions Chemical Equations The following chemical equation tells us that propane gas and oxygen gas react to form carbon dioxide gas and water vapor C3 H8 ( g) + O2 (g) CO2 (g) + H2 O( g) Propane Carbon dioxid e Oxygen Water But while it tells us what the reactants and products are and the physical state of each, it is incomplete because it is not balanced Balancing Equations To balance a chemical equation – begin with atoms that appear only in one compound on the left and one on the right; in this case, begin with carbon (C) which occurs in C3H8 and CO2 C3 H8 (g) + O2 (g) 3CO2( g) + H2 O(g) – now balance hydrogens, which occur in C3H8 and H2O C3 H8 (g) + O2 (g) 3CO2( g) + 4H2 O(g) – if an atom occurs as a free element, as for example Mg or O2, balance this element last; in this case O2 C3 H8 (g) + 5O2 ( g) 3CO2 (g) + 4H2 O( g) Balancing Equations Practice problems: balance these equations Ca( OH) 2 ( s) + HCl( g) Calcium hydroxide CO2 ( g) + H2 O(l) CaCl2 (s) + H2 O( l) Calcium chlorid e photosynthesis C4 H1 0 ( g) + O2 (g) Butane C6 H1 2 O6 (aq) + O2 (g) Glucose CO2 (g) + H2 O(g) Balancing Equations Solutions to practice problems Ca( OH) 2 ( s) + 2 HCl(g) Calcium hydroxide 6CO2 (g) + 6H2 O(l) CaCl2 (s) + H2 O( l) Calcium chlorid e photosynthesis C4 H1 0 ( l) + 13 O2 (g) 2 Bu tane C6 H1 2 O6 (aq) + 6O2 (g) Glucose 4 CO2 (g) + 5 H2 O( g) – it is common practice to use only whole numbers; therefore, multiply all coefficients by 2, which gives 2C4 H1 0 ( l) + 13O2 (g) Butane 8CO2 ( g) + 10H2 O(g) Formula Weight Formula weight: the sum of the atomic weights in atomic mass units (amu) of all atoms in a compound’s formula Ionic Comp ou nds Sod ium chlorid e (N aCl) 23.0 amu N a + 35.5 amu Cl = 58.5 amu Nickel(II) ch loride h yd rate 58.7 amu N i + 2(35.5 amu Cl) + 12(1.0) amu H) + 6(12.0 amu O) = 237.7 amu (N iCl 2• 6H 2O) Molecu lar Comp ou nds Water (H 2O) Aspirin (C9H 8 O 4) 2(1.0 amu H) + 16.0 amu O = 18.0 amu 9(12.0 amu C) + 8(1.0 amu H) + 4(16.0 amu O) = 180.0 = amu Formula Weight formula weight can be used for both ionic and molecular compounds; it tells nothing about whether a compound is ionic or molecular molecular weight should be used only for molecular compounds in this text, we use formula weight for ionic compounds and molecular weight for molecular compounds The Mole Mole (mol) – a mole of the amount of substance that contains as – – – – many atoms, molecules, or ions as are in exactly 12 g of carbon-12 a mole, whether it is a mole of iron atoms, a mole of methane molecules, or a mole of sodium ions, always contains the same number of formula units the number of formula units in a mole is known as Avogadro’s number Avogadro’s number has been measured experimentally its value is 6.02214199 x 1023 formula units per mole Molar Mass Molar mass: the formula weight of a substance expressed in grams Glucose, C6H12O6 – molecular weight: 180 amu – molar mass: 180 g/mol – one mole of glucose has a mass of 180 g Urea, (NH2)2CO – molecular weight 60.0 amu – molar mass: 60.0 g/mol – one mole of urea has a mass of 60.0 g Molar Mass We can use molar mass to convert from grams to moles, and from moles to grams You are given one of these and asked to find the other Grams of A Moles of A Use molar mass (g/mol) as the conversion factor – calculate the number of moles of water in 36.0 g water 1 mol H2 O 36.0 g H 2O x = 2.00 mol H2 O 18.0 g H 2O Grams to Moles Calculate the number of moles of sodium ions, Na+, in 5.63 g of sodium sulfate, Na2SO4 – first we find the how many moles of sodium sulfate – the formula weight of Na2SO4 is 2(23.0) + 32.1 + 4(16.0) = 142.1 amu – therefore, 1 mol of Na2SO4 = 142.1 g Na2SO4 5.63 g Na2 SO4 x 1 mol Na2SO4 = 0.0396 mol Na2SO4 142.1 g Na2 SO4 – the formula Na2SO4 tells us there are two moles of Na+ ions per mole of Na2SO4 + 0.0396 mol Na2 SO 4 2 mol Na x 1 mol Na2SO 4 = 0.0792 mol Na+ Grams to Molecules A tablet of aspirin, C9H8O4, contains 0.360 g of aspirin. How many aspirin molecules is this? – first we find how many mol of aspirin are in 0.360 g 0.360 g aspirin x 1 mol aspirin 180.0 g aspirin = 0.00200 mol aspirin – each mole of aspirin contains 6.02 x 1023 molecules – the number of molecules of aspirin in the tablet is 0.00200 mole x 6.02 x 1023 molecules = 1.20 x 1021 molecules mole Stoichiometry Stoichiometry: the study of mass relationships in chemical reactions – following is an overview of the the types of calculations we study You are given one of these Grams of A And asked to find one of these Moles of A From grams to moles, use molar mass (g/mol) as a conversion factor Moles of B Grams of B From moles to moles, From moles to grams, use the coefficients in use molar mass (g/mol) the balanced equation as a conversion factor as a conversion factor Stoichiometry Problem: how many grams of nitrogen, N2, are required to produce 7.50 g of ammonia, NH3 N2 (g) + 3H2 (g) 2NH3 ( g) – first find how many moles of NH3 are in 7.50 g of NH3 7.50 g NH 3 x 1 mol NH 3 17.0 g NH 3 = mol NH 3 – next find how many moles of N2 are required to produce this many moles of NH3 7.50 g NH 3 x 1 mol NH 3 17.0 g NH 3 x 1 mol N2 2 mol NH 3 = mol N 2 Stoichiometry Practice problem (cont’d) – finally convert moles of N2 to grams of N2 and now do the math 7.50 g NH 3 x 1 mol NH 3 17.0 g NH 3 x 1 mol N2 2 mol NH 3 x 28.0 g N 2 1 mol N2 = 6.18 g N 2 Stoichiometry Practice problems: – what mass of aluminum oxide is required to prepare 27 g of aluminum? Al2 O3 ( s) electrolysis Al( s) + O2 ( g) – how many grams each of CO2 and NH3 are produced from 0.83 mol of urea? ( NH2 ) 2 CO(aq) + H2 O Urea urease 2NH3 (aq) + CO2 (g) Limiting Reagent Limiting reagent: the reagent that is used up first in a chemical reaction – consider this reaction of N2 and O2 N2 (g) + O2 (g) before reaction (moles) 5.0 after reaction (moles) 4.0 1.0 0 2 NO( g) 0 2.0 – in this experiment, there is only enough O2 to react with 1.0 mole of N2 – O2 is used up first; it the limiting reagent – 4.0 moles of N2 remain unreacted Limiting Reagent Practice Problem – suppose 12 g of carbon is mixed with 64 g of oxygen and the following reaction takes place C(s) + O2 ( g) CO2 ( g) – complete the following table. Which is the limiting reagent? C before reaction (g) 12 g before reaction (mol) after reaction (mol) after reaction (g) + O2 CO2 64 g 0 Percent Yield Actual yield: the mass of product formed in a chemical reaction Theoretical yield: the mass of product that should be formed according to the stoichiometry of the balanced chemical equation Percent yield: actual yield divided by theoretical yield times 100 Percent yield = Actual yield x 100 Theoretical yield Percent Yield Practice problem: – suppose we react 32.0 g of methanol with excess carbon monoxide and get 58.7 g of acetic acid – complete this table CH3 OH + CO before reaction (g) 32.0 before reaction (mol) th eoretical yield (mol) th eoretical yield (g) actual yield (g) percent yield (%) excess CH3 COOH 0 58.7 Reactions Between Ions Ionic compounds, also called salts, consist of both positive and negative ions When an ionic compound dissolves in water, it dissociates to aqueous ions NaCl(s) H2 O Na+ (aq) + Cl- (aq) What happens when we mix aqueous solutions of two different ionic compounds? – if two of the ions combine to form a water-insoluble compound, a precipitate will form – otherwise no physical change will be observed Reactions Between Ions Example: – suppose we prepare these two aqueous solutions Solution 1 Solution 2 AgNO3 (s) NaCl(s) H2 O H2 O Ag+ ( aq) + NO3 - (aq) + Na ( aq) + Cl ( aq) – if we then mix the two solutions, we have four ions present; of these, Ag+ and Cl- react to form AgCl(s) which precipitates + + Ag (aq) + NO3 ( aq) + Na (aq) + Cl ( aq) + AgCl(s) + Na (aq) + NO3 ( aq) Reactions Between Ions – we can simplify the equation for the formation of AgCl by omitting all ions that do not participate in the reaction Net ionic equation: Ag+ ( aq) + Cl-(aq) AgCl( s) – the simplified equation is called a net ionic equation; it shows only the ions that react – ions that do not participate in a reaction are called spectator ions Reactions Between Ions In general, ions in solution react with each other when one of the following can happen – two of them form a compound that is insoluble in water – two of them react to form a gas that escapes from the reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid + HCO3 ( aq) + H3 O (aq) CO2 ( g) + 2 H2 O( l) Bicarbonate ion Carbon d ioxide – an acid neutralizes a base (Chapter 8) – one of the ions can oxidize another (Section 4.7) Reactions Between Ions Following are some generalizations about which ionic solids are soluble in water and which are insoluble – all compounds containing Na+, K+, and NH4+ are soluble in water – all nitrates (NO3-) and acetates (CH3COO-) are soluble in water – most chlorides (Cl-) and sulfates (SO42-) are soluble; exceptions are AgCl, BaSO4, and PbSO4 – most carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and hydroxides (OH-) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH4OH which are soluble in water Oxidation-Reduction Oxidation: the loss of electrons Reduction: the gain of electrons Oxidation-reduction (redox) reaction: any reaction in which electrons are transferred from one species to another Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a solution of copper(II) sulfate – some of the zinc metal dissolves – some of the copper ions deposit on the zinc metal – the blue color of Cu2+ ions gradually disappears In this oxidation-reduction reaction – zinc metal loses electrons to copper ions 2+ Zn(s) Zn (aq) + 2 e Zn is oxidized – copper ions gain electrons from the zinc 2+ Cu ( aq) + 2 e - Cu( s) Cu 2+ is red uced Oxidation-Reduction – we summarize these oxidation-reduction relationships in this way electrons flow from Zn to Cu2 + 2+ Zn(s) + Cu (aq) loses electrons ; gains electrons ; is red uced is oxidized gives electrons tak es electrons to Cu 2+ ; is th e from Zn; is th e oxidizin g agent red ucing agent 2+ Zn ( aq) + Cu( s) Oxidation-Reduction Although the definitions of oxidation (loss of electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others – for example, the combustion of methane CH4 (g) + O2 ( g) CO2 (g) + H2 O( g) Methane An alternative definition of oxidation-reduction is – oxidation: the gain of oxygen or loss of hydrogen – reduction: the loss of oxygen or gain of hydrogen Oxidation-Reduction – using these alternative definitions for the combustion of methane electrons are trans ferred from carb on to oxygen CH4 (g) gain s O and los es H; is oxidized + O2 (g) gain s H; is reduced is the reducin g is th e oxid izing agent agent CO2 (g) + H2 O(g) Oxidation-Reduction Five important types of redox reactions – combustion: burning in air. The products of complete combustion of carbon compounds are CO2 and H2O. – respiration: the process by which living organisms use O2 to oxidize carbon-containing compounds to produce CO2 and H2O. The importance of these reaction is not the CO2 produced, but the energy released. – rusting: the oxidation of iron to a mixture of iron oxides 4Fe(s) + 3O2 ( g) 2Fe2 O3 ( s) – bleaching: the oxidation of colored compounds to products which are colorless – batteries: in most cases, the reaction taking place in a battery is a redox-reaction Heat of Reaction In almost all chemical reactions, heat is either given off or absorbed – example: the combustion (oxidation) of carbon liberates 94.0 kcal per mole of carbon oxidized C( s) + O2 (g) CO2 (g) + 94.0 kcal/mole C Heat of reaction: the heat given off or absorbed in a chemical reaction – exothermic reaction: one that gives off heat – endothermic reaction: one that absorbs heat – heat of combustion: the heat given off in a combustion reaction; all combustion reactions are exothermic Chemical Reactions End Chapter 6