Chapter 13 Chapter 13: Chemical Reactions

1

Chapter 13

ChemicalReactions

Chapter 13: Chemical Reactions

Homework: All questions on the “Multiple-Choice” and the odd-numbered questions on“Exercises” sections at the end of the chapter.

Copyright© Houghton Mifflin Company. All rights reserved. 13 | 2

Properties

• Properties – the characteristics of asubstance

• Physical Properties – do not describe thechemical reactivity of the substance– Density, hardness, phase, color, melting point,

electrical conductivity, specific heat

• Chemical Properties – reflect the ways inwhich a substance can be transformed intoanother substance– Describe a substance’s chemical reactivity

– Burning, rusting, decomposition

Section 13.1


Physical and Chemical Propertiesand Changes

Section 13.1 Copyright© Houghton Mifflin Company. All rights reserved. 13 | 4

Chemical Reactions

• Chemical Reaction – a change that alters thechemical composition of a substance andresults in the formation of one or more newsubstances

– Decomposition of water H2O into hydrogen (H2)and oxygen (O2) gases

• A chemical reaction is simply arearrangement of the atoms. Some of theoriginal chemical bonds are broken and newbonds form.

– New and different chemical structures are formed.

Section 13.1


Chemical Reaction – Rearrangement ofAtoms


Chemical Reactions

• Generally, the atom’s valence electrons arethe only ones directly involved in the chemicalreaction.

– The nucleus is unchanged and therefore theidentity of the atom is unaffected.

• Consider the following generalized reaction

• A + B C + D

• Reactants – original substances A +B

• Products – the new substances C + D

• “” means “reacts to form” or “yields”

Section 13.1

2


Common Symbols in Chemical Equations


Chemical Reactions

• In any chemical reaction three thingsoccur:

1) Reactants disappear or are diminished.

2) New substances are formed as products.

• These products have different chemical/physical properties from the originalreactants.

3) Energy is either released or absorbed.

• Heat, light, electricity, sound

Section 13.1


Chemical Equations

• A chemical equation can be written forevery chemical reaction.

• The correct chemical formulas must beused and the equation must bebalanced.

– A balanced equation reflects the correctchemical formulas and the correct ratios ofreactants and products.


Balancing Chemical Reactions

• Most chemical reactions can bebalanced by trial and error, using threesimple principles:1) An equal number of atoms of each kind

must be represented on each side of thereaction arrow.

2) The formulas may not be changed, onlythe coefficients in front of the formulas.

3) The final set of coefficients used shouldbe the smallest whole numbers that willsatisfy the equation.

Section 13.1


Balancing Chemical Reactions –an Example

• HI H2 + I2 -- an unbalanced equation

• The formulas cannot be changed but acoefficient of “2” can be used, as below.

• 2 HI H2 + I2 -- a balance equation

• The following equations are alsoinappropriate:– HI ½ H2 + ½ I2 Fractions should not be

used

– 4 HI 2H2 + 2I2 Smallest whole numbershould be used


Tips to Balance Equations

• You must be able to properly count theatoms.– 4Al2(SO4)3 -- in this equation there are 8 Al

atoms, 12 S atoms, and 48 O atoms

• Start by balancing an element that ispresent in only one place on both sides ofthe reaction.– In the reaction C + SO2 CS2 + CO start by

balancing the S or the O.

• Insert the lowest coefficient possible oneither side to get the same number of atomsof that element on each side.

Section 13.1

3


Tips to Balance Equations

• When polyatomic ions remain intact duringthe reaction, balance them as a unit.

• If both sides balance only by the use of afractional coefficient in one place, multiply allthe coefficients by the denominator of thefraction.

– C2H2 + 5/2 O2 2 CO2 + H2O

– Multiply by 2 (the denominator of 5/2)

– 2 C2H2 + 5 O2 4 CO2 + 2 H2O

– Resulting in a balance equation with no fractions


Balancing Equations – An Example

• Mg + O2 MgO – an unbalancedequation

• Balance the oxygens first.

• Mg + O2 2 MgO– Oxygens are balanced but not

magnesiums.

• 2 Mg + O2 2 MgO– Now both magnesiums and oxygens are

balanced.

Section 13.1


Balancing EquationsConfidence Exercise

• NaN3(s) Na(s) + N2(g)

– This is an unbalance reaction.

• 2 NaN3(s) Na(s) + 3 N2(g)

– Nitrogens are now balanced.(6 on each side)

• 2 NaN3(s) 2 Na(s) + 3 N2(g)

– Both nitrogens and sodiums are balanced.

• This reaction shows how car air bags inflateby the electrical ignition of sodium azide(NaN3) to produce nitrogen gas (N2)


Combination Reactions

• Combination Reaction – at least tworeactants combine to form one product

– A + B AB

Section 13.1


Decomposition Reactions

• Decomposition Reaction – only one reactantis present and it breaks into two, or more,products

– AB A + B


Chemical Reactions – A Change in Energy

• During a chemical reaction, some chemicalbonds are broken and other bonds are formed.

• The different bonding energies of the reactantsand products result in a change in energy.

• In order to break bonds, energy must beabsorbed.

• When new bonds are formed, energy isreleased.

• Energy from a chemical reaction is released orabsorbed in the form of light, heat, electricalenergy, or sound.

Section 13.2

4


Exothermic Reactions

• Exothermic reaction – a chemicalreaction that results in a net release ofenergy to the surroundings

• Example: The burning of methane

• CH4 + 2 O2 CO2 + 2 H2O + energy

• Bonding energy in the products is lessthan bonding energy in the reactants

– therefore, energy is released.


Exothermic Reaction• Net release of energy (ER) - the bonds in the products have less

total energy than the bonds in the reactants

Section 13.2


Endothermic Reactions

• Endothermic reaction - a chemicalreaction that results in a net absorptionof energy from the surroundings

• Example: 3 O2 + energy 2 O3

• Bonding energy in the reactants is lessthan bonding energy in the products.– Although energy is released when the

ozone bonds are formed, the amount isless than is absorbed in breaking theoxygen molecule bonds

– therefore energy is absorbed.


Endothermic Reaction• Net absorption of energy (ER) - the bonds in the reactants have

less total energy than the bonds in the products

Section 13.2


Activation Energy

• Activation energy (Eact) – the energynecessary to start a chemical reaction

• In order to burn methane, one must providean initial spark to break the “first” C—H andO—O bonds.

– After the initial bonds are broken, the energyreleased breaks the bonds of still more CH4 andO2, continuously giving off energy as heat andlight.

• Eact – the minimum kinetic energy at collidingmolecules must possess to react chemically


The activationenergy required fora common match isacquired throughfriction (heat.)

Section 13.2

5


Heat Flow in Exothermic andEndothermic Reactions

Exothermic – vessel heats up asheat flows to surroundings

Endothermic – vessel cools off asheat flows in from surroundings


Explosive & Combustive Reactions

• Explosion – occurs when an exothermicchemical reaction liberates energyalmost instantaneously

• Combustion reaction – a substancereacts with oxygen by bursting intoflames and forming an oxide(exothermic reaction)

– Burning of natural gas, coal, paper, wood

Section 13.2


Hydrocarbon Combustion

• All C-H compounds (hydrocarbons) andC-H-O compounds produce energywhen they react with oxygen.

– Exothermic chemical reactions

– Give off CO2 and H2O when combustion iscomplete


Complete Hydrocarbon CombustionAn Example

• One of the components of gasoline is thehydrocarbon named heptane, C7H16. Writethe balanced equation for its completecombustion.

• Write heptane plus oxygen w/ reaction arrow

• C7H16 + O2

• The products are always CO2 & H2O.

• C7H16 + O2 CO2 + H2O

• Balance equation

• C7H16 + 11 O2 7 CO2 + 8 H2O

Section 13.2


Complete Hydrocarbon CombustionConfidence Exercise

• Write and balance the equation for thecomplete combustion of the hydrocarbonpropane, C3H8 , a common fuel gas.

• Write propane plus oxygen w/ reaction arrow

• C3H8 + O2

• The products are always CO2 & H2O

• C3H8 + O2 CO2 + H2O

• Balance equation

• C3H8 + 5 O2 3 CO2 + 4 H2O


Incomplete Combustion of Heptane

• With insufficient time or oxygen heptane doesnot completely combust, resulting in thefollowing chemical reaction:

• C7H16 + 9 O2 4 CO2 + 2 CO + C + 8 H2O

• In many cases this incomplete combustioncan be seen in the dark exhaust gases givenoff by some automobiles.

– Note that sooty black carbon (C) and poisonouscarbon monoxide (CO) are also products ofincomplete hydrocarbon combustion.

Section 13.2

6


Rate of Reaction

• How fast a reaction proceeds (rate ofreaction) depends on four variables :

1) Temperature of the reactants

2) Concentration of the reactants

3) Surface Area of the reactants

4) Catalyst Presence


Temperature affects Rate of Reaction

• Recall that temperature is the average kineticenergy of the molecules.

• For chemical reactions to proceed, thereacting molecules must collide with enoughkinetic energy to break bonds. (Eact)

• Increased kinetic energy (higher temperature)of the reacting molecules results in moreabundant and harder molecular collisions,therefore, the rate of reaction increases.

Section 13.2


Concentration affects Rate of Reaction

• Generally, if the concentration of thereactants is greater, the rate of thereaction is also greater

• A higher concentration of reactantsresults in more molecular collisions anda faster reaction rate.

• In pure oxygen (100%) a normalcigarette will actually burst into flames.– Recall that typical air has only about 21%

oxygen.


Concentration affects Rate of ReactionPhosphorus burning in pure oxygen (left) vs. 21% oxygen (right)

Section 13.2


Surface Area affects Rate of Reaction

• As the surface area of reactants increase, sodoes the reaction rates.

• For example, chemical weathering of rocksprogresses much faster as the surface areaof the rocks increase.

• Dramatic and devastating explosions haveoccurred when large amounts of coal dust orgrain dust are present.

– It is the high surface area of these dusts thatcauses combustion with explosive speed.


Catalysts affects the Rate of Reaction

• Catalyst – a substance that increases the rateof reaction, but is not consumed in thereaction

• Some catalysts provide a surface to aid inconcentrating the reactants.

• Most catalysts provide a new reactionpathway that has a lower activation energy.

• Although not consumed, the catalyst usuallyforms an intermediate “product” that takespart in the process and then decomposesback to its original form.

Section 13.2

7


Catalysts

• Generally provides a new reaction pathway with a lower activationenergy requirement. The presence of NO lowers the activationenergy requirement (Ecat)


Catalysts affects the Rate of ReactionAn Example

• In the manufacture of H2SO4 the reaction isvery slow without a catalyst.

• 2 SO2 + O2 2 SO3 (slow)

• When NO is added two fast reactions occur.

• 2 NO + O2 2 NO2 (fast)

• 2 SO2 + 2 NO2 2 SO3 + 2 NO (fast)

• Therefore the following is the net reaction:

• 2 SO2 + O2 2 SO3 net reaction (fast)

Section 13.2


Catalysts affects the Rate of ReactionAnother Example

• The decomposition of H2O2 at roomtemperature is very slow.

• 2 H2O2 2 H2O + O2 (slow)

• When a small amount MnO2 is added to theH2O2 the reaction proceeds rapidly.

– Catalyst is signified by placing it over the reactionarrow.

Section 13.2

MnO2• 2 H2O2 2 H2O + O2


Automotive Catalytic Converter

• Beads of Pt, Rh, or Pd serve as catalysts to quicklyconvert noxious CO and NO into CO2 and N2.

Section 13.2


Enzymes

• In the biological world, living organismsalso use catalysts. They are calledenzymes.

• These enzymes control variousphysiologic reactions.

• For example, milk sugar (lactose) isbroken down in a reaction catalyzed bythe enzyme lactase.

• Individuals who are “lactose intolerant”have a deficiency of lactase.


Bombardier BeetleThis exothermic reaction uses an enzyme catalyst.

Section 13.2

8


Acids and Bases

• Early in the history of chemistry,substances were classified as acids orbases

• One of the first theories to explain acidsand bases was put forward in 1887 bythe Swedish chemist, Svante Arrhenius

• He proposed that the properties ofaqueous solutions of acids and basesare due to the hydrogen ion (H+) andthe hydroxide ion (OH-), respectively


Acid

• When dissolved in water, acids have thefollowing properties:

– Conducts electricity

– Litmus dye blue to red

– Sour taste (never taste an acid!)

– Reacts with and neutralizes a base

– Reacts with active metals, liberating H gas

Section 13.3


Base

• When dissolved in water, bases havethe following properties:

– Conduct electricity

– Litmus dye red to blue

– Reacts with and neutralizes an acid


An Arrhenius Acid

• When the colorless, gaseous substancehydrogen chloride (HCl) is added towater, virtually all the HCl moleculesionize into H+ ions and Cl- ions

• The acidic properties of an acid are dueto the hydrogen ions (H+)

• In the case of HCl, virtually all of theatoms ionize into H+ ions and Cl- ions

Section 13.3


An Arrhenius Acid

• As the HCl molecule disassociates inwater the hydrogen ions are transferredfrom the HCl to the water molecules, asshown in the following:

• HCl + H2O H3O+ + Cl-

– The H3O+ is called the hydronium ion

• Therefore, an Arrhenius acid is asubstance that produces hydrogen ions,H+, (or hydronium ions, H3O

+) in water


HCl Reacts with H2O forming -Hydronium ions (H3O

+) and Chlorine ions (Cl-)

Section 13.3

9


Strong & Weak Acids

• An acid classified as strong, ionizes almostcompletely in solution

• Strong acids include HCl, HNO3, H2SO4

– All three of these in water ionize virtuallycompletely

• Weak acids do not ionize to any great extentin solution– At equilibrium, only a small fraction of the

molecules disassociate to form H3O+

• Acetic acid, HC2H3O2 (vinegar) is a commonweak acid


Strong Acids Ionizemore Completely

• ConductElectricity Better

• Bulb GlowsBrightly

Section 13.3


Dynamic Equilibrium

• A double arrow may be used if thereverse reaction is significant

• When two ‘competing’ reactions areoccurring at the same time, the systemis in dynamic equilibrium

• In the case above, only a small fractionof the acetic acid moleculesdisassociate

Section 13.3

• HC2H3O2 + H2O H3O+ + C2H3O2


Acids are Very Useful Compounds

• Sulfuric acid has a number of differentindustrial uses

• Refining petroleum, processing steel,fertilizers

• Dilute HCl helps digest food in the stomach

• Citric acid (H3C6H5O7) in citrus fruits

• Carbonic (H2CO3) and Phosphoric (H3PO4)acids in soft drinks

• Acetic acid (HC2H3O2) in vinegar

Section 13.3


An Arrhenius Base

• When pure sodium hydroxide (NaOH) isadded to water, it dissolves releasingNa+ and OH-

• The basic properties of NaOH are dueto the hydroxide ions (OH-)

• Therefore an Arrhnius base is asubstance that produces hydroxide ions(OH-)


An Arrhenius Base

• Ammonia (NH3) is also considered abase even though it does not contain(OH-)

• In water ammonia reacts with H2O toform (OH-)

• Common household bases include Drano(NaOH) and Windex (NH3)

Section 13.3

• NH3 + H2O NH4+ + OH

10


Water

• Water will slightly ionize by itself

• Therefore all aqueous solutions have both thehydronium ion (H3O

+) and the hydroxide ion(OH-)

• For pure water the concentrations of H3O+ and

OH- are equal, and therefore neutral

• An acidic solution has more H3O+

• A basic solution has more OH-

Section 13.3

• H2O + H2O H3O+ + OH-


Acidic and Basic Solutions

Acidic Solution Basic Solution

Acidic solution – higher concentration of H3O+

Basic solution – higher concentration of OH-

Section 13.3


pH – “power of hydrogen”

• The relative acidity or basicity of asolution is commonly designated byciting the pH

• The pH of a solution is a logarithmicmeasure of the concentration of thehydrogen ion

• Since the pH scale is logarithmic, eachdrop/rise of 1 on the scale is actually atenfold increase/decrease in acidity


The pH Scale

A solution with a pH of 7 is neutral, a solution with apH less than 7 is acidic, and a solution with a pHgreater than 7 is basic

Section 13.3


Body Fluids and pH

• Most body fluids of a healthy personmust remain in a very narrow range onthe pH scale

• Thus the pH of different body fluids maybe used as a diagnostic measure

– The pH of blood should be between 7.35-7.45


Acid-Base Reaction

• When an acid is brought in contact with abase its characteristic properties disappear

– And vice versa

• This is known as an acid-base reaction

• An acid and a hydroxide base react toproduce water and a salt

• A salt is an ionic compound composed of anycation except H+ and any anion except OH-

– Examples of salt include, KCl, Ca3(PO4)2,CuSO4

.5H2O, CaSO4.2H2O

Section 13.3

11


CuSO4.5H2O is blue –

anhydrous CuSO4 is white

Adding water will once again hydrate the substance


White Sand National Monument, NMThe Sand is composed of the salt CaSO4

.2H2O

Copyright © Bobby H. Bammel. All rights reserved.

Section 13.3


Acid-Base Reaction – An Example

• Write a balanced equation for stomach acidHCl and milk of magnesia Mg(OH)2

• HCl is an acid and Mg(OH)2 is a base

• \ we know the water and a salt is produced

• HCl + Mg(OH)2 H2O + a salt

• Determine the salt from the given reactants

• HCl + Mg(OH)2 H2O + MgCl2• Balance the equation, using Mg2+

• 2 HCl + Mg(OH)2 2 H2O + MgCl2


Acid-Base Reaction – Confidence Exercise

• Write a balanced equation for stomach acidHCl and aluminum hydroxide (Di-Gel) Al(OH)3

• HCl is an acid and Al(OH)3 is a base

• \ we know the water and a salt is produced

• HCl + Al(OH)3 H2O + a salt


• HCl + Al(OH)3 H2O + AlCl3• Balance the equation, using Al3+

• 3 HCl + Al(OH)3 3 H2O + AlCl3

Section 13.3


Acid-Carbonate Reaction

• An acid and a carbonate or hydrogencarbonate react to give CO2, H2O, and asalt

• For example the antacid Tums containsCaCO3 and can also be used toneutralize too much stomach acid


Acid-Carbonate Reaction – An Example

• Write the balance equation for the reactionbetween stomach acid (HCl) and Tums(CaCO3)

• This is an acid-carbonate reaction, \ theproducts are CO2, H2O, and a salt

• HCl + CaCO3 CO2 + H2O + a salt


• HCl + CaCO3 CO2 + H2O + CaCl2• Balance the equation, using Ca2+

• 2 HCl + CaCO3 CO2 + H2O + CaCl2Section 13.3

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Acid-CarbonateReaction

• 2HCl+CaCO3CO2+H2O+CaCl2


Acid-Carbonate Reaction –Confidence Exercise

• Write the balance equation for the reactionbetween nitric acid (HNO3) and Na2CO3

• This is an acid-carbonate reaction, \ theproducts are CO2, H2O, and a salt

• HNO3 + Na2CO3 CO2 + H2O + a salt


• HNO3 + Na2CO3 CO2 + H2O + NaNO3

• Balance the equation

• 2 HNO3 + Na2CO3 CO2 + H2O + 2 NaNO3

Section 13.3


Double-Replacement Reactions

• Both acid-base and acid-carbonate aretypes of double-replacement reactions

• In this type of reaction the positive andnegative components “change partners”

• The general format is as follows …

• AB + CD AD + CB


Double-Replacement Reactions

Have the format AB + CD AD + CB, basically allfour components (K, Cl, Ag, NO3) “change partners”

AB + CD AD + CB

Section 13.3


Double-Replacement Precipitation Reaction

• When two aqueous solutions of salts aremixed, a precipitate often results

• One pair of ionsform a precipitate,the other pairremains indissolved


Double-Replacement Precipitation Reaction- An Example

• Write the equation for the double-replacementreaction between KI (aq) and Pb(NO3)2 (aq)

• KI (aq) + Pb(NO3)2 (aq) AD + CB

• 2 KI (aq) + Pb(NO3)2

(aq) 2 K(NO3) (aq) +PbI2 (s)

Section 13.3

13


Double-Replacement PrecipitationReaction - Confidence Exercise

• Write the equation for the double-replacement reaction between Na2SO4

(aq) and BaCl2 (aq)

• Na2SO4 (aq) + BaCl2 (aq) AD + CB

• Complete the equation by exchangingpartners and balancing the equation

• Na2SO4 (aq) + BaCl2 (aq)2 NaCl (aq) +BaSO4 (s)


Oxidation & Reduction

• Oxidation- oxygen combines with a substance– Or when an atom or ion loses electrons

• Reduction – oxygen removed from asubstance– Or when an atom or ion gains electrons

• All of the electrons lost by an atom or ion mustbe gained by other atoms or ions– therefore oxidation and reduction occur at the same

time.

• These type of reactions are called oxidation-reduction reactions or redox reactions forshort.

Section 13.4


Oxidation-Reduction ReactionSteel-Making Reduction of Hematite Ore

• Hematite (Fe2O3) must be reduced with coke(C) in a furnace to for pure iron.

• 2 Fe2O3 + 3 C 4 Fe + 3 CO2

• Fe3+ Fe

– Each Fe gains 3 electrons – reduction

• C C4+

– Each C loses 4 electrons - oxidation

• Therefore hematite (iron oxide) has beenreduced


Activity

• The relative activity of any metal is itstendency to lose electrons to ions ofanother metal or to hydrogen ions.

• A metal’s activity can be determined byplacing the metal in a solution thatcontains ions of another metal.

• If the test metal replaces the metal insolution, then the test metal is said to bemore active.

Section 13.4


Activity Series

• Shows therelativeactivity of anumber ofcommonmetals


Single-Replacement Reaction

• If an element a is listed above element B inthe activity series …

– A is more active than B.

– A will replace B in a aqueous solution.

• This is called a single-replacement reaction.

• The general format for this type of reaction is

• A + BC B + AC

• For example Zn and CuSO4

Section 13.4

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Single-Replacement Reactions

• A + BC B + AC

– Cu precipitates, but the ZnSO4 is soluble.

– Zn loses electrons \ has been oxidized

– Cu gains electrons \ has been reduced


Zinc replaces copper ionstherefore Zn is more active than Cu.

Zn + CuSO4 Cu + ZnSO4

Section 13.4


Single-Replacement ReactionAn Example

• Refer to the activity series table and predictwhether placing copper metal in a solution ofsilver nitrate will lead to a reaction. If so,complete and balance the equation.

• Cu is above Ag, therefore a reaction willoccur.

• Cu will the oxidized to Cu2+.

• Ag+ will be reduced to Ag.

• NO3- is a polyatomic ion and will stay intact.


Single-Replacement ReactionAn Example (cont.)

• The unbalanced equation for the reaction is

• Cu + AgNO3(aq) Ag + Cu(NO3)2(aq)

• Balance the equation.

– Remember that the NO3- stays intact and is

balanced as a unit.

• Cu + 2 AgNO3(aq) 2 Ag + Cu(NO3)2(aq)

• Cu loses electrons. (oxidation)

• Ag gains electrons. (reduction)

Section 13.4


Single-Replacement Reaction BetweenCopper Wire and Silver Nitrate

Cu + 2 AgNO3(aq) 2 Ag + Cu(NO3)2(aq)


Single-Replacement ReactionConfidence Exercise

• Suppose you place Al metal in a solutionof CuSO4 (aq), and also put a strip of Cumetal in a solution of Al2(SO4)3(aq).

• Refer to the activity series and predictwhich single-replacement reaction willtake place.

• Complete and balance the equation forthe reaction.

Section 13.4

15


Activity Series

• Note that Al ismore active thanCu, therefore Alwill replace theCu in CuSO4.


Single-Replacement ReactionConfidence Exercise (cont.)

• Al + CuSO4(aq) Cu + Al2(SO4)3(aq)

• Balance the equation.– Remember that SO4

2- stays intact and isbalanced as a unit.

• 2 Al + 3 CuSO4(aq) 3 Cu +Al2(SO4)3(aq)

• Al loses electrons. (oxidation)

• Cu gains electrons. (reduction)

Section 13.4


Single-Replacement Reaction with Acids

• Metals above hydrogen in the activityseries will undergo a single-replacementreaction with acids, giving off hydrogengas and a salt of the metal.


Single-Replacement Reaction with Acid

• Fe reacts with H2SO4 to give off H and an Fesalt

Section 13.4


Single-Replacement Reactions with Water

• Metals above magnesium in the activityseries will undergo a vigorous single-replacement reaction with liquid water,giving off hydrogen gas and the metalhydroxide.

• For example Ca reacts with wateraccording to the following equation:

• Ca + 2 H2O H2 + Ca(OH)2

– Ca loses electrons. (oxidation)

– H gains electrons. (reduction)


Active Metals and WaterMetals more active than Mg will react with water.

Section 13.4

16


A Summary of Reaction Types


ReactionTypes

Section 13.4


A Mole & Avogadro’s Number

• Mole – the quantity of a substance thatcontains as many formula units as there areatoms in exactly 12 grams of 12C

– Mole is abbreviated “mol.”

• In exactly 12 grams of 12C there are exactly6.023 x 1023 12C atoms.

• This huge number (6.023 x 1023) is referred toas Avagodro’s Number.

– The Italian physicist Avogadro was the first personto use the term molecule.

• Therefore a mole is 6.023 x 1023 units.


A Mole & Formula Mass

• A mole of any substance has a massequal to the same number of grams asits formula mass. (FM)

• For example: Copper has a FM of63.5m– Therefore a mole of Cu contains 6.023 x

1023 atoms and has a mass of 63.5 grams.

• Water has a FM of 18.0 m– A mole of H2O contains 6.023 x 1023

molecules and has a mass of 18 grams.

Section 13.5


One Mole of Six Substances63.5 g of Cu, 27.0 g of Al, 55.8 g of Fe, 32.1 g of S,

253.8 g of I, and 200.6 g of Hg


Avogadro’s Number

• 6.023 x 1023 is a huge number

• If we poured out 6.023 x 1023 BB’s onthe U.S., they would cover the entirecountry to a depth of 4 inches!

• But how do we actually know how manyparticles there are in a mole?

Section 13.5

17


A Mole = 6.023 x 1023 units

• We know that it takes 96,485 coulombs (C) toreduce 1 mole of singly charged ions toatoms.– For example Na+ + e- Na

• Thus, 96,485 C must be the total charge onone mole of electrons.

• Single electron has a charge of 1.6022 x 10-19

C.

• Therefore Avogadro’s number must be

• (96,485 C/mol)/(1.6022 x 10-19 C/electron)

• = 6.023 x 1023 electrons/mole

Section 13.5

Chapter 13 Chapter 13: Chemical Reactions

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Transcript of Chapter 13 Chapter 13: Chemical Reactions