Chemistry C2 Topic 4
Revision PowerPoint
4
What I’m Looking For
1) To be able to explain the meaning of
the term rate of reaction.
2) To be able to use data, equations and
graphs to calculate reaction rates.
Making reactions go faster
This PowerPoint covers Topic 3 Lessons 18 to 22
C2_15 Key words 1
C2_15 Key words 2
Collision theory
Rate of reaction
Chemical reactions occur when reactant molecules collide with
enough energy for chemical bonds to break – these are called
successful collisions
The rate of a chemical reaction is a measure of how quickly the
reactants are converted into the reaction products.
Rate of reaction = Amount of product made ÷ time
Rate of reaction = Amount of reactant used ÷ time
Kinetic theory
All matter (solids, liquids and gases) is composed of particles.
The speeds of those particles and how closely packed together
the particles are depends on the temperature of their
surroundings.
Rate curves
Rate curves show the progress of the reaction graphically. The
gradient of the curve shows how quickly or slowly the reaction
is going
WILF 1
WILF 2
Some chemical reactions take a long time to
happen. The rusting of iron is a very slow
chemical reaction.
An explosion is a chemical reaction which has an
incredibly fast reaction rate.
The rate of reaction can be calculated from
o the amount of product made ÷ the time
taken, or,
o the amount of reactant used ÷ the time
taken
WILF 2
Region(1)
Thegradientissteepest.Thereaction
isfastestinregion1.
Rate(1)=20cm3 /60s…0.34cm3/s
Region(3)
Thegradientiszero.The
reactionhasnow
stopped.
Rate(2)=0cm3 /30s
…0cm3/s
Region(2)
Thegradientislesssteep.
Thereactionisnow
slowingdown.
Rate(2)=10cm3 /120s
…0.08cm3/s
Thistableofresultsshowsthatafter30
secondsofreacting,15cm3 ofgaswas
produced.
So,after30secondstherateofreaction
=15cm3 ÷ 30s
=0.5cm3/s
What I’m Looking For
1) To be able to recall that chemical
reactions happen when particles
collide.
2) To be able to describe and explain the
effect of changing temperature on
the rate of a reaction.
C2_19 Key words 1
C2_19 Key words 2
More frequent collisions
Activation energy
Molecules in a warm solution collide more frequently than
the same molecules in a cold solution. Many molecules are
moving faster in the warm solution because of an increase
in kinetic energy that the extra thermal energy gives them.
This is the minimum amount
of energy needed to start a
chemical reaction.
It is also the minimum amount
of energy the colliding
particles need to react.
The activation energy can be
thought of as an ‘energy
barrier’.
Successful collisions
Successful collisions result in the chemical bonds in the
molecules breaking. Increasing the temperature of
the chemicals can therefore result in many more
successful (violent) collisions.
WILF 1
Chemical reactions happen when there are
frequent successful collisions between molecules
of reactants.
The chances of successful collisions happening
increases if:
o the reactant particles move faster and hit
each other harder (more violently)
o the number of particles of reactant in a
space is increased (i.e. if the chemicals are
more concentrated)
WILF 2
In the picture above, two
hydrogen molecules are colliding
with one oxygen molecule.
If the molecules are travelling
fast enough and hit each other
very hard, then there is enough
energy to overcome the energy
barrier for the reaction. When
this happens, chemical bonds
break and new molecules (water)
can then form. Gentle collision
do not lead to a chemical
reaction.
Whenstudentscarriedoutthereactionbetween
sodiumthiosulphate(25cm3 of0.2Msolution)
withhydrochloricacid(5cm3 of1Msolution)at
highertemperaturestheyfoundthatthetime
takenforthe‘X’todisappearfromview
decreased.Warmersolutions=fasterreactions
Whenthetemperatureofthesolutionsis
increased,themoleculesofreactantsmove
faster,collidewitheachothermoreoftenand
moreviolently.Therearemoresuccessful
collisionsbecausemoremoleculeshavemore
thantheminimumenergyneededto
overcomethereaction’senergybarrier.
What I’m Looking For
1) Use particle collision theory to
explain the effect of changing
concentration on the rate of reaction
2) Use particle collision theory to
explain the effect of changing the
pressure of gases on the rate of
reaction .
C2_20 Key words
The pressure of a gas is linked to the number of
collisions between the gas molecules and the
container they are trapped inside. Pressure is
measured in pascals.
The concentration of a solution depends upon
the mass of dissolved solute, in grams, present
in a decimetre cubed of solution. (i.e. grams
per decimetre cubed,
g/dm3).
WILFS 1
WILF 1
Chemical reactions happen when there
are frequent successful collisions
between molecules of reactants.
Atahigherconcentration,therearemoresolute
molecules inthesamespace(volume).Therewill
bemoresuccessfulcollisionspersecondbetween
thereactantmolecules.
The chances of more of these
successful collisions happening per
second increases if the solutions
reacting have a higher concentration of
solute.
WILF 1 Extension
WILF 2
Whena2cmstripof
magnesiumribbonreacts
withhydrochloricacid
wecantimehowlongit
takesforthemagnesium
ribbontoreactandform
asolution.
Ifwedothisforsolutions
ofincreasing
concentration thetime
forthemagnesiumto
completelyreact
decreasesasshownin
thegraphontheright.
Chemical reactions happen when there are
frequent successful collisions between
molecules of reactants.
The chances of more of these successful
collisions happening per second increases if
the gases reacting together are at a higher
pressure
WILF 2
Atahigherpressure,
therearemoregas
moleculesinthesame
space(volume)and
thereforetherewillbe
moresuccessful
collisionspersecond
betweenthereactant
molecules.
Numberofsuccessful
collisionsbetweenacid
moleculesandmagnesium
atomsincreasesfromleft
toright
Higherconcentration=shorterreactiontime
What I’m Looking For
1) To be able to describe and explain the
effects of changing the surface area on
the rate of a reaction.
2) Use the particle-energy model for
chemical reactions and the idea of
‘contact between reactant molecules’
when explaining these effects.
Key words and their meanings
Powders haveaverylargesurfaceareatovolume
ratio,i.e.moreofthesolidisontheoutside.Lumps
havesmallsurfaceareatovolumeratios.Mostofthe
solidliesunderneaththesurfaceandcannotbe
reached.
Toachieveafastreactionitisbetterforsolidstobe
crushedintopowderssothereisalargercontactarea
betweenthesolidandtheliquid.
Crushedmarblepowderreactsmuchfasterwith
hydrochloricacidthanifitwasinlumps.Thesolid
hasmuchlargersurfacearea ifitiscrushedintoa
powder.
WILF 1
Twoseparateexperiments
involving5gofmarbleand15cm3
of1.0mol/dm3 hydrochloricacid
werecarriedout.Inthefirst
experimentthe5gofmarblewas
crushedtoafinepowder,inthe
secondthe5gofmarblewasleftas
largechips.
Thegraphoftheresultsshows
thattherateofreactioninvolving
powderedmarblewasmuchfaster,
asshownbythesteepnessofthe
gradientofthegraphline.The
reactiontook6minutesto
completewithpowder,but8
minuteswithmarblechips.
WILF 1
Solidscanbefoundaslargepieces,smalllumpsandpowders.In
chemicalreactionsinvolvingsolids.itisalwaysbesttousea
powderedsolid.Theyreactfasterthanlargepieces.
Powderedsolidhasalargesurface
area.Thepowderwillhavemuch
bettercontactwiththeliquiditis
reactingwiththanifyouusedthe
samemassbutinlumps.
Theliquidcanonlyreactwiththe
outsideofthelump.Whenthe
lumpiscrushed,moleculesonthe
insideofthelumpbecome
exposedonthesurfaceafter
crushing.
Theinsideofthelumpdoes
notmeettheliquidtobegin
with,andwillneedtowait
fortheouterpartsofthe
solidtoreactfirst.
WILF 2
Theexperimentshowsthatthepowderreactedfasterthanthe
lumps.Thisisbecausethepowderhasamuchlargersurface
areaandthereforemorepowderwasabletobeincontact
withhydrochloricacidmolecules.
Withpowderthereweremoremarble-acidmolecular
collisionspersecond.Thepowderreactiontookashorted
timetocomplete.
What I’m Looking For
1) To be able to explain what
catalysts are and what they do.
2) To be able to explain why
catalysts are important in
industry.
Key words and their meanings
oCatalyticconvertersarepartofcarexhaustsystemsand
thecatalystsareusedtoquicklymakeharmfulgases
fromtheengine,likecarbonmonoxide,oxidesof
nitrogenandunburnedfuel,safe.
oAcatalyst isachemicalthatspeedsuptherateofaslow
reaction.
oThe‘energybarrier’forachemicalreactionistheenergy
requiredtogetthereactionstarted.Wecallthisthe
activationenergy.
WILF 1
WILF 2
Theresultsoftheexperimentusing2g,
0.2gor0gofcatalystinexperimentsshow
thatcatalystshelpproducemoreoxygen
thanthesameexperimentwithcatalyst.
2gofcatalysthelpstomakeanevenfaster
reactionthan0.2gofcatalyst,itisnot10
timesbetter!!
Catalystsarechemicalsubstancesthatcanspeedup
therateofachemicalreactionwithoutbeingusedup.
Catalystsarenorreactants.2gofmanganesedioxide
catalystweighedoutatthestartofareactioninvolving
hydrogenperoxide,willleave2gattheendofthe
reaction.
Catalystsworkbyinterferingwiththechemicalbonds
inthereactantssothatevengentlecollisionsbetween
reactantsresultinthemoleculesreacting.Catalysts
aresaidtolowertheactivationenergyforthereaction.
WILF 2
Catalystsworkbyweakening
thechemicalbondsbetween
theatomsinthemoleculesof
reactants.
Theenergybarrierforthe
reactionisnowlower.
Nowgentlecollisionsbetween
reactantmoleculescannow
causechemicalbondstobreak.
Thecatalysthasloweredthe
activationenergy forthe
reaction,asshowninthe
diagram.
Withoutacatalystthereisa
highactivationenergy
Withacatalystbeingused
theactivationenergyislower
WILF 2
Catalystsareimportantinindustrybecause:
• catalystsspeedupveryslowreactions
• thesamemassofcatalystcanbeusedoverandover
again
• catalystshelptoreducethecostofproducing
chemicals
Examplesoftheuseofcatalystsinindustryinclude:
• ironfilingstohelpnitrogenandhydrogenreact
fasterintheproductionofammonia
• nickelintheconversionofliquidplantoilsinto
solidmargarine
• aluminiumoxidebeadsinhelpingtocrack
hydrocarbonsintheoilindustry