Chemical Kinetics Intro (Rates of Reaction) Unit 3: Kinetics and Equilibrium
What actually happens during a chemical reaction? • Reactant molecules must collide for a reaction to occur
Chemical Kinetics • The study of rates of chemical reactions and the mechanisms (or steps) by which a chemical reaction takes place. • Reaction rates vary greatly – some are very fast (burning) and some are very slow (disintegration of a plastic bottle in sunlight).
Definition: • Rate of reaction: the change in concentration of a reactant or product per unit of time. • Rxn Rate (avg) = Δ[reactant or product] Δ time • Note: [square brackets] = mol/L = concentration
Example : If the concentration of reactant X after 50 s in a reaction was 0.0079 mol/L and after 100 s was 0.0065 mol/L, calculate the average reaction rate. • Rate = Δ[X] Δt = (0.0065 – 0.0079 M) (100-50 s) = - 0.000028 M/s
Think about it! • Concentration of reactant decreases with time, as it is being used up • The rate is fastest when concentration of reactants is greatest; and slows when concentration of reactants are less
Still thinking….. • The reactants are being used up as the reaction takes place. • What’s happening to the concentration of products over time? • It is increasing.
Concentration rate of rxn • In the beginning of the reaction, product is formed quickly • But slows over time because less reactants means less reactants collisions producing the new product. Concentration of product (mol/L) Time (s)
Factors that affect the Rate: • Concentration (and pressure) • Temperature • Amount of surface area • Catalysts • Nature of reactants
Concentration • A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate • Note: We are not increasing the amount being made for a given balanced equation with limiting reactants, we are only speeding up how quickly those products are made.
Pressure (Gases) • Affects the rate of reaction, particularly gases. • When you increase the pressure, the molecules have less space in which they can move. That greater concentration of molecules increases the number of collisions.
Temperature • Remember: Temperature is a measure of the average kinetic energy of a system • So higher temperature implies higher average kinetic energy of molecules and more collisions per unit time.
Surface Area • Reducing the size of particles increases the rate of a reaction because it increases the surface area available for collisions to take place. This increases the number of collisions.
Catalysts • A catalyst is a substance that speeds up a reaction without being used up itself. • Some reactions have catalysts that can speed them up, but for many reactions there is no catalyst that works. • How do they work? (We will talk about this later!)
Nature of Reactants • Reactants with large number of chemical bonds that have to be broken, or with strong chemical bonds that have to be broken, will lead to a slower rate. (Later….. • The strong bonds will cause a high activation energy, leading fewer collisions being successful. • More complex reactant molecules with more atoms, will have more difficulty lining up properly to have a successful reaction. This will also lead to a lower reaction rate.)
Some good rules of thumb… • Ions will react more quickly, because already broken up. • Simple molecules (non-metals covalently bonded) will take more time • Complex molecules and ionic compounds (metal and non-metal bonded) will take the most time because the strong ionic bonds need lots of energy to be broken, and the complex molecules will be hard to have proper geometry (line up) • Summary (fast to slower): Ions (solutions)>simple molecular>ionic and complex molecular
Try it : • Place the following reactions from fastest to slowest, based on the previous slides. • 2Cu+2 + 2I- 2 Cu+ + I2 • 2NaCl 2Na + Cl2 • Cl2 2Cl • Page 467 #1, 3, 4, 5
Answer from previous slide: • 1 is fastest, then 3, and 2 is slowest