Rates of Reaction Flashcards
rate of reaction
the rate of a chemical reaction is defined as the change in concentration per unit time of any one reactant or product
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
theory behind experiment
hydrogen peroxide decomposes slowly into water and oxygen gas, the rate of decomposition can be greatly increased with the addition of the catalyst, manganese dioxide
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
equation
H2O2 -MnO2-> H2O + 1/2 O2
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
where do you place the manganese oxide?
in a weighing bottle in the hydrogen peroxide solution
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
first step
knock the weighing bottle into the hydrogen peroxide using the cotton thread and stop clock started immediately
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
where do you collect the oxygen gas
from the inverted graduated cylinder (downward displacement of water)
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
recording
volume of oxygen recorded every 3 seconds until it is constant
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
how does manganese dioxide look
black powder
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
repeat
without catalyst and results compared
MONITORING THE ARTE OF PRODUCTION OF OXYGEN FROM HYDROGEN PEROXIDE, USING MANGANESE DIOXIDE AS A CATALYST
conclusion
the rate of the reaction at the start of the experiment was very fast, however as time went on the rate began to slow down and eventually stopped (level graph)
how to measure average rate
total volume of oxygen/total time
to find instantaneous rate
draw a tangent to the cur
find the slope of the tangent
5 factors that affect the rate of a chemical reaction
nature of chemicals (ionic/covalent) particle size (one is solid) concentration temperature catalysts
which are faster in general ionic or covalent reactions
ionic
why are ionic reactions usually faster
coming together of ions vs bonds formed again
the larger the size of the particle
the slower the reaction - less surface area
if finely divided particles are used
a dust explosion may happen
5 conditions necessary for a dust explosion to happen
dust must be: combustible, dry
oxygen present
enclosed space
source of ignition
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
reactants
calcium carbonate (marble) and dilute hydrochloric acid
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
equation
CaCO3 + 2HCl -> CaCl2 + H2O + CO2
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
first step
weigh out conical flask and cotton wool
weigh marble chips and add to flask with cotton in mouth
weigh put dilute HCl and quicly add to flask (remove plug and put back)
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
after HCl added
quickly put flask on electronic balance and start the stop clock
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
stop clock has started
take mass every 30 seconds
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
what is the loss in mass due to
the carbon dioxide lost through the cotton wool plug
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
how to calculate loss in mass
subtracting each mass from the initial mass before reaction started
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
graph
loss in mass over time
HOW PARTICLE SIZE CAN AFFECT REACTION RATE
how to prove for particle size
graph results for different sizes of marble chips
the greater the concentration
the faster the reaction, generally
why is it faster for greater concentrations
a greater chance of reactants colliding with each other successfully
if temperature increased, and why
faster, more kinetic energy, more collisions
do catalysts always speed up reaction rates
no
catalyst that slows down the decomposition of hydrogen peroxide
glycerine
if a catalyst slows something down what is it called
a negative catalyst or inhibitor
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
equation
Na2S2O3 + 2HCl -> s(↓) + 2NaCl + SO2
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
what can you measure
the visibility of the cross drawn underneath
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
first step
known volume and molarity of sodium thiosulfate poured into a conical flask, placed on a piece of paper with an x drawn on it
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
step 2
known volume and molarity of HCl added and stop clock started
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
step 3
swirl flask and stop clock when cross no longer visible
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
how to examine concentrations
procedure repeated using other concentrations of sodium thiosulfate and times notes
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
graph for concentrations
rate vs M
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
how to do for temperature
note temperature at the start and then repeat using various temperatures, heated by bunsen burne r
TO STUDY THE EFFECTS ON THE REACTION OF (I) CHANGING CONCENTRATION (II) CHANGING TEMPERATURE, USING SODIUM THIOSULFATE AND HYDROCHLORIC ACID
graph for temperature
reaction vs absolute temperature
catalyst
a substance that alters the rate of a chemical reaction but is not used up during the reaction
4 general properties of catalysts
remain chemically unchanged
specific
reversible in their action
catalyst poisons
may catalysts be physically changed by the end of a reaction?
yes, just not chemically
what does it mean to be reversible in their action
they catalyse both forward and reverse reactions to the same extent, do not have an effect on the position of equilibrium, but you do get to equilibrium faster
3 types of catalysts
heterogenous catalysts
homogenous catalysts
autocatalysts
heteorgenous catalysts
when the reactants and the catalyst are in different phases e.g in haber process the reactants are gases and catalyst, iron is a solid
homogenous
when the reactants and the catalyst are in the same phase e.g potassium iodide soln (liquid) catalyses decomposition of hydrogen peroxided (liquid) into water and oxygen
autocatalysts
when the catalyst is the product of the reaction
catalysts in catalytic converters
platinum, palladium and rhodium
reactions that occur in catalytic converters
conversion of harmful gases such as carbon monoxide and nitrogen monoxide to form carbon dioxide and nitrogen
activation energy
the minimum amount of energy that colliding particles must have in order for a reaction to occur
size of activation energy determines what?
the rate of the reaction
can you increase or decrease activation energy for a reaction
no
if activation energy is high
only small amount of molecules have it so reaction is slow
what happens if you increase the temperature 2
increase in the kinetic energy of the molecules
greater proportion of molecules will have activation energy
why is increase in kinetic energy not hugely significant?
the molecules will collide more often, but if they do not have energy the collisions will be ineffective
2 theories for the mechanisms of catalysts
surface absorption theory
intermediate compound formation theory
what does surface absorption theory explain
heterogenous catalysis
according to surface absorption theory, if a solid catalyst is placed among gaseous reactants, what happens
the gases adsorb onto the surface of the catalyst
surface absorption theory
advantages of gases adsorbing onto catalyst
increased concentration of the gases on the surface of the catalyst, more effective collisions happening
surface absorption theory
what happens to the product
leaves the surface of the catalyst and allows more reactants to adsorb
example of heterogenous catalysis
the haber process
intermediate compound formation theory
2 points about the intermediate compound
forms very quickly and decomposes as soon as it is formed
to get evidence for intermediate compound formation theory
potassium sodium tartrate is oxidised by hydrogen peroxide using cobalt (II) as a catalyst
first few steps of experiment to get evidence for intermediate compound formation theory
potassium sodium tartrate dissolved in water, cobalt (II) ions added, hydrogen peroxide added
colour of Co2+ ions
pink
to get evidence for intermediate compound formation theory
colour when everything added
green
to get evidence for intermediate compound formation theory
what is given ff
carbon dioxide and steam given off very vigorously
to get evidence for intermediate compound formation theory
after a while when the reaction is finished
pink colour appears again, Co2+ not used up
effect of catalysts on activation energy
when a catalyst is added to a reaction it provides an alternative route with a lower activation energy, more molecules will have this, rate increases
