应用化学专业英语lesson10ChemicalEquilibriumandkinetics
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Temperature can also have an effect. For exothermic reactions reactants products + heat
Raising the temperature shifts it to the left. For endothermic reactions heat + reactants products
Dynamic Equilibrium
5 - 24
Chemical equilibrium
Dynamic process Rate of forward Rxn = Rate of reverse Rxn H2O(l) (reactant) H2O(g) (product)
Concentration of reactants and products remain constant over time.
At Equilibium
2SO3(g)
SO2(g)+O2(g)
Initially
SO3(g)
Initially
5 - 28
Figure 9.10
N2(g) + O2(g)
At Equilibium
2NO(g)
N2(g)+O2(g)
Initially
NO(g)
Initially
5 - 29
Equilibrium
effects of catalysts and enzymes. how to control a reaction.
5-4
Reaction Rates
Speed at which reactant is used up. Speed at which product forms.
Fast:
Oxidation: Paper burning
Lesson 10 Chemical Equilibrium and Kinetics
5-1
Reaction Rates
2H2(g) + O2(g)
2H2O + Energy
H2(g) + O2(g) may
Energy
-H
Very stable product stay together for (H < 0) lifetime without reacting to form
Enthalpy & Entropy
Kinetics
Only tell us if a reaction will occur but not how long it will take.
Measures the time required for a reaction to occur.
5-3
Chemical kinetics
NH3
Add more NH3?
Reaction shifts to the left [N2] and [H2] inc
5 - 35
Le Chatelier’s principle
Adding Pressure affects an equilibrium with gases
N2(g) + 3 H2(g)
Lower Eact More collisions Uncatalysed reaction Catalysed reaction
Lower activation energy
Alters reaction mechanism but not products Is not used up during the reaction.
Reaction rates can be affected by : 1. Reactant structure(polar vs. nonpolar) • physical state of reactants (vapor vs liq.) 2. Concentration of reactants (medications) • surface area (sugar cube vs crystals) 3. Temperature (hypothermia & metabolism) 4. Catalyst (H2O2 & blood)
Kinetics of a chemical reaction can tell us -
how long it will take for a reaction to reach completion. how chemicals react to form products (mechanism).
Keq =
[C]c [D]d [A]a [B]b
moles per liter
5 - 31
Figure 9.11
Equilibrium constant (K)
aA + bB reactants cC + dD products
Keq =
[C]c [D]d [A]a [B]b
5 - 32
Equilibrium constant (K)
Rxn Progress
wBaidu Nhomakorabeater.
Just because something has the potential to react doesn’t mean it will do so immediately.
5-2
Chemical kinetics
The study of reaction rates (speed)
2 NH3(g)
N2
Add more N2?
Reaction shifts to the right [NH3] inc, [H2] dec
5 - 34
Le Chatelier’s principle
Adding or removing reagent
N2(g) + 3 H2(g)
2 NH3(g)
Dynamic Equilibrium
5 - 25
H2O(l) (reactant)
H2O(g)
Equilibrium and reaction rates
Reaction rate
(product) A point is ultimately reached where the rates of the forward and reverse reactions are the same.
2. They have to be aligned correctly. 3. They have to have enough E.
(Parked cars don’t collide)
5-8
Activation Energy
The activation energy Eact Is the minimum energy needed for a reaction to take place upon proper collision of reactants.
Kinetic Region Concentration Equilibrium Region
Concentration of reactants and products remain constant over time.
Time
5 - 30
Equilibrium constant (K)
Equilibrium expression (for any reaction at constant temperature) aA + bB reactants cC + dD products coefficients
5 - 14
Reaction Rates
Catalyst: 3. Adding a Catalyst:
Lower Eact More collisions
Uncatalysed reaction
5 - 15
Reaction Rates
Catalyst: 3. Adding a Catalyst:
N2 O2 N2 O2
Molecules are not aligned correctly.
5-7
Effective collisions
For reactants to make products: 1. Molecules must collide
(solvents really help)
5 - 11
Reaction Rates
Concentration : 1. More Reactants:
More cars More collisions If Increase reactant concentration then Increase # of collisions so Increase reaction rate.
4 mol of reactants
2 NH3(g)
Add P?
2 mol of products
Increasing pressure causes the equilibrium to shift to the side with the least moles of gas.
5 - 36
Le Chatelier’s principle
5 - 16
Reaction Rates
Catalyst: 3. Adding a Catalyst:
Lower Eact More collisions Uncatalysed reaction Catalysed reaction
Lower activation energy
Enzymes are biological catalysts.
Temperature: 2. Higher Temperature:
Faster cars More collisions
More Energy More collisions
Reacting molecules move faster, providing colliding molecules w/ Eact.
5-9
Energy diagrams
A temporary state where bonds are reforming.
Activated Complex
Show the E during a reaction.
Energy
Activation energy
Eact
-H
5 - 10
Factors Influencing Rxn Rates
5 - 12
Reaction Rates
Concentration: 1. More Reactants:
More surface area More collisions
8 blocks: 34 surfaces 8 blocks: 24 surfaces
5 - 13
Reaction Rates
5 - 17
Catalytic Converter
5 - 18
Catalytic Converter
5 - 19
Equilibrium
A state where the forward and reverse conditions occur at the same rate.
I’m in static equilibrium.
At this point, equilibrium is achieved. Time
5 - 26
Figure 9.8
2SO2(g) + O2(g)
At Equilibium
2SO3(g)
SO2(g)+O2(g)
Initially
SO3(g)
Initially
5 - 27
Figure 9.9
2SO2(g) + O2(g)
Slow:
Oxidation: Nails rusting Paper turning yellow
5-5
Figure 9.1
Reaction Rates Fast:
Slow:
Slower:
5-6
Effective collisions
A reaction won’t happen if: Insufficient energy to break bonds.
N2(g) + 3 H2(g)
Keq =
[ NH3 ] 2 [ N2 ] [ H2 ] 3
2 NH3(g)
5 - 33
Le Chatelier’s principle
Stress causes shift in equilibrium Adding or removing reagent
N2(g) + 3 H2(g)
Raising the temperature shifts it to the left. For endothermic reactions heat + reactants products
Dynamic Equilibrium
5 - 24
Chemical equilibrium
Dynamic process Rate of forward Rxn = Rate of reverse Rxn H2O(l) (reactant) H2O(g) (product)
Concentration of reactants and products remain constant over time.
At Equilibium
2SO3(g)
SO2(g)+O2(g)
Initially
SO3(g)
Initially
5 - 28
Figure 9.10
N2(g) + O2(g)
At Equilibium
2NO(g)
N2(g)+O2(g)
Initially
NO(g)
Initially
5 - 29
Equilibrium
effects of catalysts and enzymes. how to control a reaction.
5-4
Reaction Rates
Speed at which reactant is used up. Speed at which product forms.
Fast:
Oxidation: Paper burning
Lesson 10 Chemical Equilibrium and Kinetics
5-1
Reaction Rates
2H2(g) + O2(g)
2H2O + Energy
H2(g) + O2(g) may
Energy
-H
Very stable product stay together for (H < 0) lifetime without reacting to form
Enthalpy & Entropy
Kinetics
Only tell us if a reaction will occur but not how long it will take.
Measures the time required for a reaction to occur.
5-3
Chemical kinetics
NH3
Add more NH3?
Reaction shifts to the left [N2] and [H2] inc
5 - 35
Le Chatelier’s principle
Adding Pressure affects an equilibrium with gases
N2(g) + 3 H2(g)
Lower Eact More collisions Uncatalysed reaction Catalysed reaction
Lower activation energy
Alters reaction mechanism but not products Is not used up during the reaction.
Reaction rates can be affected by : 1. Reactant structure(polar vs. nonpolar) • physical state of reactants (vapor vs liq.) 2. Concentration of reactants (medications) • surface area (sugar cube vs crystals) 3. Temperature (hypothermia & metabolism) 4. Catalyst (H2O2 & blood)
Kinetics of a chemical reaction can tell us -
how long it will take for a reaction to reach completion. how chemicals react to form products (mechanism).
Keq =
[C]c [D]d [A]a [B]b
moles per liter
5 - 31
Figure 9.11
Equilibrium constant (K)
aA + bB reactants cC + dD products
Keq =
[C]c [D]d [A]a [B]b
5 - 32
Equilibrium constant (K)
Rxn Progress
wBaidu Nhomakorabeater.
Just because something has the potential to react doesn’t mean it will do so immediately.
5-2
Chemical kinetics
The study of reaction rates (speed)
2 NH3(g)
N2
Add more N2?
Reaction shifts to the right [NH3] inc, [H2] dec
5 - 34
Le Chatelier’s principle
Adding or removing reagent
N2(g) + 3 H2(g)
2 NH3(g)
Dynamic Equilibrium
5 - 25
H2O(l) (reactant)
H2O(g)
Equilibrium and reaction rates
Reaction rate
(product) A point is ultimately reached where the rates of the forward and reverse reactions are the same.
2. They have to be aligned correctly. 3. They have to have enough E.
(Parked cars don’t collide)
5-8
Activation Energy
The activation energy Eact Is the minimum energy needed for a reaction to take place upon proper collision of reactants.
Kinetic Region Concentration Equilibrium Region
Concentration of reactants and products remain constant over time.
Time
5 - 30
Equilibrium constant (K)
Equilibrium expression (for any reaction at constant temperature) aA + bB reactants cC + dD products coefficients
5 - 14
Reaction Rates
Catalyst: 3. Adding a Catalyst:
Lower Eact More collisions
Uncatalysed reaction
5 - 15
Reaction Rates
Catalyst: 3. Adding a Catalyst:
N2 O2 N2 O2
Molecules are not aligned correctly.
5-7
Effective collisions
For reactants to make products: 1. Molecules must collide
(solvents really help)
5 - 11
Reaction Rates
Concentration : 1. More Reactants:
More cars More collisions If Increase reactant concentration then Increase # of collisions so Increase reaction rate.
4 mol of reactants
2 NH3(g)
Add P?
2 mol of products
Increasing pressure causes the equilibrium to shift to the side with the least moles of gas.
5 - 36
Le Chatelier’s principle
5 - 16
Reaction Rates
Catalyst: 3. Adding a Catalyst:
Lower Eact More collisions Uncatalysed reaction Catalysed reaction
Lower activation energy
Enzymes are biological catalysts.
Temperature: 2. Higher Temperature:
Faster cars More collisions
More Energy More collisions
Reacting molecules move faster, providing colliding molecules w/ Eact.
5-9
Energy diagrams
A temporary state where bonds are reforming.
Activated Complex
Show the E during a reaction.
Energy
Activation energy
Eact
-H
5 - 10
Factors Influencing Rxn Rates
5 - 12
Reaction Rates
Concentration: 1. More Reactants:
More surface area More collisions
8 blocks: 34 surfaces 8 blocks: 24 surfaces
5 - 13
Reaction Rates
5 - 17
Catalytic Converter
5 - 18
Catalytic Converter
5 - 19
Equilibrium
A state where the forward and reverse conditions occur at the same rate.
I’m in static equilibrium.
At this point, equilibrium is achieved. Time
5 - 26
Figure 9.8
2SO2(g) + O2(g)
At Equilibium
2SO3(g)
SO2(g)+O2(g)
Initially
SO3(g)
Initially
5 - 27
Figure 9.9
2SO2(g) + O2(g)
Slow:
Oxidation: Nails rusting Paper turning yellow
5-5
Figure 9.1
Reaction Rates Fast:
Slow:
Slower:
5-6
Effective collisions
A reaction won’t happen if: Insufficient energy to break bonds.
N2(g) + 3 H2(g)
Keq =
[ NH3 ] 2 [ N2 ] [ H2 ] 3
2 NH3(g)
5 - 33
Le Chatelier’s principle
Stress causes shift in equilibrium Adding or removing reagent
N2(g) + 3 H2(g)