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OCR A A-Level Inorganic Chemistry > Chapter 18 > Flashcards

Chapter 18 Flashcards

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1
Q

How is rate of reaction calculated? What is its units?

A
  • Change in concentration/ change in time
  • Its units are therefore mol dm^-3 s^-1
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2
Q

How is rate of reaction linked to the concentration of reactants?

A
  • The rate of reaction is directly proportional to the concentration of a reactant raised to a power
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3
Q

What is this power known as?

A
  • The order of reaction (for that reactant, as different reactants can have different orders)
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4
Q

List 3 types of orders, and what each one means for the rate of reaction.

A
  • Zero order; the concentration doesn’t affect the rate
  • First order; the rate of reaction is directly proportional to the change in the change in concentration (if the concentration of the reactant doubles, the rate doubles)
  • Second order; the change in the rate of reaction is directly proportional to the square of the change in concentration of the reactant (if the concentration of the reactant doubles, the rate of reaction increases by a factor of 4)
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5
Q

What value is used for the rate of reaction when calculating the order of a reaction? How is it found?

A
  • The initial rate (at t=0)
  • From doing a reaction and plotting the change in concentration over time
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6
Q

How can you find the change in concentration of a reactant?

A
  • Titrations
  • Colorimetry
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7
Q

What do the rate-concentration graphs of zero, first and second order reactions look like?

A
  • Zero order: straight horizontal line
  • First order: linear line, crosses at the origin
  • Second order: positive quadratic, turning point at the origin
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8
Q

What is the rate equation?

A
  • rate = k [A]^m [B]^n
  • k= rate constant
  • A and B: reactants
  • m: order of reaction with respect to A
  • n: order of reaction with respect to B
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9
Q

What is the rate constant? What does it tell you?

A
  • The number that connects the rate of reaction and the concentrations (and orders)
  • The higher k is, the higher the rate of reaction
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10
Q

What is the overall order of a reaction? How would this be calculated from the rate equation?

A
  • The sum of the orders in respect to each reactant
  • m+n
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11
Q

Explain the shape of rate-concentration graphs for zero-order reactions.

A
  • Straight horizontal line since rate=k
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12
Q

Explain the shape of rate-concentration graphs for first-order reactions.

A
  • Linear line through the origin since rate=k[A]
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13
Q

Explain the shape of rate-concentration graphs for second-order reactions.

A
  • Positive quadratic through the origin since rate=k[A]^2
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14
Q

What is continuous monitoring? What is it used for?

A
  • Continuous measurements taken during the course of the reaction
  • To plot concentration-time graphs
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15
Q

What are three properties that can be monitored for the purpose of plotting a concentration-time graph?

A
  • Gas produced
  • Mass loss
  • Colour change
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16
Q

What device is used to measure colour change?

A
  • A colorimeter
17
Q

How do colorimeters work?

A
  • Since solutions transmit light that is the same colour as them and absorb all of the other colours, colorimeters pass light of the complementary colour (colour opposite in the colour wheel) to the solution through the solution
  • The amount of light absorbed is measured, which is directly proportional to concentration
18
Q

How can a concentration-time graph be created using a colorimeter?

A
  • Prepare standard solutions of the coloured chemical in the reaction
  • Select the filter with the complementary colour
  • Zero the calorimeter with water
  • Take readings of the absorbance of the standard solutions
  • Plot a calibration curve (concentration against absorbance)
  • Do the reaction, taking samples at timed intervals
  • Find the absorbance of these samples, and use the calibration curve to find their concentration
  • Then plot a graph of concentration against time
19
Q

What are 2 useful features of a concentration-time graph?

A
  • Its gradient is the rate of reaction
  • Its shape tells you the order with respect to a reactant for zero and first order reactions
20
Q

What is the shape of a zero order reaction on a concentration-time graph? What does its gradient show?

A
  • Straight line with a negative gradient
  • The rate constant
21
Q

What is the shape of a first order reaction on a concentration-time graph?

A
  • Exponential the shape of 1/x, but touching the y-axis when t=0
22
Q

What is special about a first order reaction? What is this pattern known as?

A
  • The time taken for the concentration of the reactant to half is constant, and is therefore known as the half-life (t1/2)
  • Exponential decay
23
Q

How can the rate constant for a first order reaction be calculated from a concentration-time graph?

A
  • ln(2) / t1/2
  • Use a concentration-time graph to find the rate of reaction
  • Divide the rate by the concentration at that time
24
Q

What is a more efficient way of finding the instantaneous rate of reaction (rate of reaction when t=0)? What is beneficial about it?

A
  • A clock reaction
  • You only take a single measurement of when a visual change happens, but you can assume that the average rate of reaction is the same as the initial rate
25
Q

What calculation is done from the results of a clock reaction, and what does the calculation tell you?

A
  • You divide 1 by the time taken for the visual change to happen
  • The initial rate is proportional to this value
26
Q

How does an iodine clock reaction work?

A
  • Hydrogen peroxide is reacted with iodide ions, which forms iodine, which turns the reaction mixture brown
  • A small volume of thiosulfate ions and starch are added to the mixture
  • Thiosulfate ions react with iodine and turn it back to iodide ions, turning the solution colourless
  • When the thiosulfate ions are used up, iodine is no longer converted into iodide ions, so it reacts with starch to turn the mixture blue-black
  • You time how long it takes for the reaction mixture to turn blue after adding the thiosulfate ions- this tells you how long it took for iodine to be formed, and therefore you can calculate the order in respect to hydrogen peroxide
  • You repeat this and change the concentration of hydrogen peroxide
27
Q

Give the equations of the reactions that happen in iodine clock reactions.

A
  • H2O2(aq) + 2H+(aq) + 2I-(aq) → I2(aq) + 2H2O(I)
  • 2S2O32-(aq) + I2(aq) → 2I-(aq) + S4O62-(aq)
28
Q

How do you calculate the order in respect to hydrogen peroxide?

A
  • You divide 1 by the time taken for the colour change to happen
  • You plot this value against the concentration of hydrogen peroxide
  • This gives you a rate-concentration graph
29
Q

What limitation is there to clock reactions?

A
  • The rate of reaction has to be measured over a short period of time, so that the calculated value is close to the true value
  • Otherwise there’s a larger change in rate
30
Q

How do reactions actually happen, and why?

A
  • They happen in multiple steps
  • It is unlikely that all of the correct particles will collide at the same time with the correct energy and orientation
31
Q

Which step in reactions is the most important, and why?

A
  • The slowest one
  • It is the rate-determining step
32
Q

How are rate equations and the rate-determining step linked?

A
  • The rate equation only includes species involved in the rate determining step (so zero-order reactants aren’t in the rate-determining step)
  • The orders in the rate equation are the same as the number of moles in the rate-determining step
33
Q

What fact are calculations of k dependent on, and why?

A
  • The temperature is constant
  • As temperature increases, the value of k increases
34
Q

Why do the rate of reaction and rate constant increase when the temperature increases?

A
  • The most dominant factor is the fact that a greater proportion of the particles have the activation energy or more
  • The particles have more kinetic energy, and therefore there is a higher frequency of collisions
35
Q

How can you find rate constants at different temperatures?

A
  • Carry out the experiment at different temperatures
36
Q

What is the Arrenhius equation? Include units.

A
  • k= A e^ (-Ea / RT)
  • k: rate constant
  • A: pre-exponential/ frequency factor
  • Ea: activation energy (J mol^-1)
  • R: gas constant
  • T: temperature in Kelvin
37
Q

Where do the 2 parts of the Arrenhius equation come from?

A
  • A takes into account the increased frequency of collisions
  • The rest of the equation represents the proportion of molecules that exceed Ea
38
Q

How can A and Ea be calculated if you have a table of k and T values?

A
  • ln(k) = -Ea / R x 1/T + ln(A)
  • y = mx +c
  • Plot a graph of ln(k) against the reciprocal of time
  • The gradient is -Ea/ R
  • The y-intercept is ln(A)

OCR A A-Level Inorganic Chemistry (11 decks)

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