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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
C 3 H 8 ( g) Propane + O 2 (g) Oxygen CO 2 (g) Carbon dioxid e + H 2 O( g) 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 C 3 H 8 and CO 2
C 3 H 8 (g) + O 2 (g) 3 CO 2 ( g) + H 2 O(g)
– now balance hydrogens, which occur in C 3 H 8 and H 2 O
C 3 H 8 (g) + O 2 (g) 3CO 2 ( g) + 4 H 2 O(g)
– if an atom occurs as a free element, as for example Mg or O 2 , balance this element last; in this case O 2
C 3 H 8 (g) + 5 O 2 ( g) 3CO 2 (g) + 4H 2 O( g)
Balancing Equations
Practice problems: balance these equations
Ca( OH) 2 ( s) Calcium hydroxide + HCl( g) CaCl 2 (s) Calcium chlorid e + H 2 O( l) CO 2 ( g) + H 2 O(l) ph otosynthes is C 6 H 1 2 O 6 (aq) Glucose + O 2 (g) C 4 H 1 0 ( g) + O 2 (g) Butane CO 2 (g) + H 2 O(g)
Balancing Equations
Solutions to practice problems
6 CO 2 Ca( OH) 2 ( s) Calcium hydroxide (g) + + 6 H 2 O(l) 2 HCl(g) ph otosynthes is CaCl 2 (s) Calcium chlorid e + H 2 O( l) C 6 H 1 2 O 6 (aq) + Glucose C 4 H 1 0 ( l) Bu tane + 13 2 O 2 (g) 4 CO 2 (g) + 5 H 2 O( g) 6 O 2 (g)
– it is common practice to use only whole numbers; therefore, multiply all coefficients by 2, which gives
2 C 4 H 1 0 ( l) Butane + 13 O 2 (g) 8 CO 2 ( g) + 10 H 2 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 (N iCl 2 • 6H 2 O) 58.7 amu N i + 2(35.5 amu Cl) + 12(1.0) amu H) + 6(12.0 amu O) = 237.7 amu Molecu lar Comp ou nds Water (H 2 O) Aspirin (C 9 H 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 10 23 formula units per mole
Molar Mass
Molar mass: the formula weight of a substance expressed in grams Glucose, C 6 H 12 O 6 – molecular weight: 180 amu – molar mass: 180 g/mol – one mole of glucose has a mass of 180 g Urea, (NH 2 ) 2 CO – 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
36.0 g H 2 O x 1 mol H 2 O 18.0 g H 2 O = 2.00 mol H 2 O
Grams to Moles
Calculate the number of moles of sodium ions, Na + , in 5.63 g of sodium sulfate, Na 2 SO 4 – first we find the how many moles of sodium sulfate – the formula weight of Na 2 SO 4 is 2(23.0) + 32.1 + 4(16.0) = 142.1 amu – therefore, 1 mol of Na 2 SO 4 = 142.1 g Na 2 SO 4 –
5.63 g Na 2 SO 4 x 1 mol Na 2 SO 4 142.1 g Na 2 SO 4
the formula Na 2 SO 4 tells us there are two moles of Na ions per mole of Na 2 SO 4
= 0.0396 mol Na 2 SO 4
+
2 mol N a + 0.0396 mol N a 2 SO 4 x = 0.0792 mol N a + 1 mol N a 2 SO 4
Grams to Molecules
A tablet of aspirin, C 9 H 8 O 4 , 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 asp irin 180.0 g asp irin = 0.00200 mol as pirin
– – each mole of aspirin contains 6.02 x 10 23 molecules the number of molecules of aspirin in the tablet is
0.00200 mole x 6.02 x 10 23 molecules mole = 1.20 x 10 21 molecules
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 And asked to find one of these Grams of A Moles of A Moles of B Grams of B From grams to moles, use molar mass (g/mol) as a conversion factor From moles to moles, use the coefficients in the balanced equation as a conversion factor From moles to grams, use molar mass (g/mol) as a conversion factor
Stoichiometry
Problem: how many grams of nitrogen, N 2 , are required to produce 7.50 g of ammonia, NH 3
N 2 (g) + 3H 2 (g) 2NH 3 ( g)
– first find how many moles of NH 3
7.50 g N H 3 x 1 mol N H 3 17.0 g N H 3
are in 7.50 g of NH
= mol NH 3
3 – next find how many moles of N 2 produce this many moles of NH 3 are required to
7.50 g N H 3 x 1 mol N H 3 17.0 g N H 3 x 1 mol N 2 2 mol N H 3 = mol N 2
Stoichiometry
Practice problem (cont’d) – finally convert moles of N 2 now do the math to grams of N 2 and
7.50 g N H 3 x 1 mol N H 3 17.0 g N H 3 x 1 mol N 2 2 mol N H 3 x 28.0 g N 2 1 mol N 2 = 6.18 g N 2
Stoichiometry
Practice problems: – what mass of aluminum oxide is required to prepare 27 g of aluminum?
Al 2 O 3 ( s) electrolysis Al( s)
+
O 2 ( g)
– how many grams each of CO 2 and NH 3 produced from 0.83 mol of urea?
ureas e ( NH 2 ) 2 CO(aq) + H 2 O 2NH 3 (aq) +
are
CO 2 (g) Urea
Limiting Reagent
Limiting reagent : the reagent that is used up first in a chemical reaction – consider this reaction of N 2
N 2 (g)
and O 2
+ O 2 (g) 2 NO( g) before reaction (moles) 5.0
1.0
0 after reaction (moles) 4.0
0 2.0
– – – in this experiment, there is only enough O 2 1.0 mole of N 2 O 2 is used up first; it the limiting reagent 4.0 moles of N 2 remain unreacted to react with
Limiting Reagent
Practice Problem – suppose 12 g of carbon is mixed with 64 g of oxygen and the following reaction takes place
C(s) + O 2 ( g) CO 2 ( g)
– complete the following table. Which is the limiting reagent?
C + O 2 CO 2 before reaction (g) before reaction (mol) 12 g 64 g 0 after reaction (mol) after reaction (g)
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 = Actu al yield Theoretical yield x 100
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
CH 3 OH + CO CH 3 COOH before reaction (g) 32.0
before reaction (mol) th eoretical yield (mol) th eoretical yield (g) actual yield (g) percent yield (%) excess 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
H 2 O NaCl(s) 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
Solu tion 1 AgNO 3 (s) H 2 O Ag + ( aq) + NO 3 (aq) Solu tion 2 NaCl(s) H 2 O 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)
Ag + (aq)
which precipitates
+ NO 3 ( aq) + Na + (aq) + Cl ( aq) AgCl(s) + Na + (aq) + NO 3 ( 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
HCO 3 ( aq) + H 3 O + (aq) CO 2 ( g) + 2 H 2 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 NH 4 + soluble in water are – all nitrates (NO 3 ) and acetates (CH 3 COO ) are soluble in water – – most chlorides (Cl ) and sulfates (SO 4 2 ) are soluble; exceptions are AgCl, BaSO 4 , and PbSO 4 most carbonates (CO 3 2 ), phosphates (PO 4 3 ), sulfides (S 2 ), and hydroxides (OH ) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH 4 OH 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 Cu 2+ ions gradually disappears In this oxidation-reduction reaction – zinc metal loses electrons to copper ions
Zn(s) Zn 2 + (aq) + 2 e Zn is oxidized
– copper ions gain electrons from the zinc
Cu 2 + ( 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 Cu 2 + Zn 2 + ( aq) + Cu( s) Zn(s) + loses electrons ; is oxidized gives electrons to Cu 2+ ; is th e red ucing agent Cu 2 + (aq) gains electrons ; is red uced tak es electrons from Zn; is th e oxidizin g agent
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
CH 4 (g) Methan e + O 2 ( g) CO 2 (g) + H 2 O( g)
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 CH 4 (g) gain s O and los es H; is oxidized + is the reducin g agent O 2 (g) gain s H; is reduced is th e oxid izing agent CO 2 (g) + H 2 O(g)
Oxidation-Reduction
Five important types of redox reactions – – – combustion: burning in air. The products of complete combustion of carbon compounds are CO 2 and H 2 O.
respiration: the process by which living organisms use O 2 CO to oxidize carbon-containing compounds to produce 2 and H 2 O. The importance of these reaction is not the CO 2 produced, but the energy released.
rusting: the oxidation of iron to a mixture of iron oxides –
4Fe(s)
bleaching:
+ 3O 2 ( g) 2Fe 2 O 3 ( s)
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) + O 2 (g) CO 2 (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