Kinetics Flashcards
What is kinetics? What is a rate?
Kinetics involves the study of the rate of a reaction (if a reaction is faster, you can do more in less time [and thus make more money]). In general a rate is a measure of the amount that something changes per unit time (how fast conc. changes with time). Reaction rate is proportional to the gradient/curve of the line. In Chem, they’re always positive.
Measure of reactant consumption (change in reactants has a minus sign, as it/they get used up and rate is always positive [rate = -Δ[R]/ΔT])/product formation (change in product concentrations divided by the time taken [rate = Δ[P]/ΔT—the brackets mean concentration]) speed. Alternative rates, slopes eventually zero.
*See doc for diagram…
The unit for rate of reaction is mol dm^-3 s^-1 (change in concentration per time). Changes in gradient show the effect of changing conditions on the rate of reaction WITHOUT NEEDING TO CONVERT THE UNITS INTO mol dm^-3 s^-1. As long as you’re using the same unit, you can compare.
Techniques to measure rate
Some techniques use the total time taken for an event to occur (average reaction rate [ex. time it takes for a piece of magnesium to disappear, time taken for the reaction to change color/cloud over, time taken to collect a certain amount of gas in a graduated cylinder, etc.]). In these cases, there’s an inverse relationship between the time for the events and the rate (ex. 1/T [see formulae in ‘What is kinetics? What is rate’ card]?).
There are many methods used in Chem to measure the rate of reaction WHILE the reaction is taking place: this is done by measuring the change in one quantity over time for a short time and is known as measuring the instantaneous rate. Common quantities to measure when measuring instantaneous rates of reaction include mass loss (of mixture), volume or pressure of gas produced (volume/pressure of gas product), change in electrical conductivity, and color change. If you measure the INSTANTANEOUS reaction rate at the very beginning of the experiment, it is called the initial reaction rate. Initial reaction rates are much more accurate than average reaction rates.
Why is the initial reaction rate more accurate than average reaction rates?
Fastest and most accurate (it’s the most linear part), as everything is known (ex. reactant temps, concentrations). Errors caused by changes in temp are minimized. Slope of line will yield [initial] reaction rate. Why it’s usually used (gradient at t = 0).
In exothermic reactions, because the surroundings increase in temp, the reaction speeds up (why measuring at the beginning [initial rate] is more accurate).
What is the pressure of a gas the result of?
The pressure of a gas (unit: kiloPascal, kPa) is the result of the collisions of gas particles with the walls of the container: any change that causes gas particles to collide more frequently with the container walls will increase the gas pressure, and vice versa (ex. high pressure in soda can, lower outside, gas wants to escape).
Why does the rate slow as the reaction progresses?
Rate is slowing because you’re using up reactants/forming more products (until reaction is complete?): reactant collisions less frequent as reaction progresses (more product molecules forming).
What is collision theory?
Collision theory dictates how we know what factors affect the reaction rate. For a reaction to occur, the reacting molecules must collide with each other (may do so gently and just bounce off each other: if they have less than the activation energy, nothing will happen) with enough kinetic energy (energy greater than the activation energy, Ea) to overcome repulsion of electron clouds and to break some bonds in the particles, and at the correct orientation (so that reactive parts come in contact with each other). The reaction will go faster or slower depending on these.
Four factors affect the rate of a reaction according to collision theory: the concentration of the solutions involved, temp, the surface area of any solids involved, and the addition of a catalyst. Concentration/Pressure and surface area mainly affect the collision rate, whereas temperature and catalyst mainly affect the proportion with the required activation energy.
What is the activation energy?
The minimum amount of energy needed to start a reaction (unit: kJmol^-1). It’s a random point (for the systems)? You can measure/calculate it: it’s a number that doesn’t really depend on anything, but it stays the same no matter what.
*See doc for enthalpy diagrams (change in enthalpy dictates the energetics part)…
You can calculate it by measuring the rate at different temps, making a 1/T graph (temp in Kelvin), with y being the natural log of the rate. The slope will be directly proportional to the gas energy (m = -Ea/R [R is the gas constant]).
Concentration (collision theory)
Remember: in order for things to react, they have to collide. The more reactant molecules in a specific volume, the greater the frequency of collision and therefore, the greater the reaction rate. For gasses, increasing the pressure is equivalent to increasing the concentration. A low concentration means that things will be spread out and they’ll find it difficult to find each other. More collisions in a given amount of time means faster reaction: the more frequently they collide, the more frequently they react. Doesn’t affect amount of particles with enough energy (ratio fixed).
Temperature (collision theory)
A higher temp means greater kinetic energy for the molecules, so greater collision frequency (particles moving quicker) and more molecules with energy GREATER (has to be greater) than the activation energy (affects for two reasons: it has a huge effect on reaction speed [why it’s better/has the greater effect]). Does not change activation energy. Generalization: for every 10° added, the rate doubles.
Surface area (collision theory)
The greater the surface area, the greater the frequency of collision and, thus, reaction rate. Dividing the solid into smaller pieces increases the surface area (when big, only a little bit to be collided with: when small, more surface area for collisions to happen). Would have to hit first layer and then, after that reacts, you have to go to the second layer. When crushed, everything is exposed, though it’s almost ineffective without stirring (helps spread everything out).
Addition of a catalyst (collision theory)
Something you put in a reaction that speeds it up/helps it happen, but goes back to where it started/isn’t used up after the reaction (doesn’t get used up: can be reused). Often metals from Sc to Znc. These are called transition metals: they can transfer electrons, and that’s what reactions are all about.
Catalysts provide an alternative route of reaction with a lower activation energy. A Maxwell-Boltzmann distribution curve (see doc) shows that if the activation energy is lowered, then many more molecules have energies greater than the activation energy, so reaction rate with a catalyst will increase (see also enthalpy graph [change in enthalpy unaffected; catalysts only affect activation energy]). You really only use a catalyst when the reaction doesn’t work at temperatures that are reasonable (speed it up). Different catalysts means different rates. Heterogeneous catalysts are not in the same state as the reactants (homogeneous are).
For biological systems, a catalyst (enzyme) can also affect the orientation of the reactants: enzymes have particular shapes (reactants can fit into the enzyme in the exact right shape, which allows the two reactants to bend/align/react [correct orientation]). Have to come together in a particular place/way (in examples, just a conceptual thing).
What is the effect of stirring?
Stirring creates more collision but doesn’t change energy.
Maxwell-Boltzmann distribution curves
Particles (in systems) have a wide distribution of energies (some have low energy, others have high). For gasses, the Maxwell-Boltzmann distribution curve shows this distribution at a given temperature.
Gives information about what proportion of particles have enough energy to react when they collide. Having energy greater than the activation energy means you have enough energy that something will happen when you collide: molecules that don’t have enough energy (nothing will happen when they collide) have to wait for the molecules with enough energy to react. Why water will evaporate when you leave it on the counter when it’s not 100℃ (some have enough energy to evaporate [when we boil, we give them enough energy]).
There’s a particular energy that most particles (in the system [?]) have: your room temp.
*See doc for curve…
Important features (Maxwell-Boltzmann distribution curve)
There are no molecules with zero energy, only a few molecules have very high energy, and there is no maximum energy for molecules.
Never touches the x-axis (would mean there’s an upper limit [hence the asymptote]: so, on the graph, you can see that when you have zero molecules, you have zero energy).
If the temperature of the sample is increased from T1 to T2 (see doc), the average kinetic energy of the molecules increases (KE ∝ T), the most probable energy of the molecules increases, the spread of energies increases, more molecules have energies greater than the activation energy, and the shape of the distribution curve changes (curve shifts [see: peak shifts to a higher energy]).
Half-range uncertainty
avg ⨦ (max - min)/2 (could calculate what percent of the avg this # is): bigger second # means less precise.
