Kinetics, Potential Energy Diagrams, and Equilibrium

Remember Heat Calculations We used these three formulas to determine how much heat a reaction absorbed or released

A Few Reminders q = mCΔT q = mHf q = mHv Used when temp. changes Temp change means NO phase change q = mCΔT q = mHf q = mHv The sign (+/-) in front of “q” tells us whether the system gained heat or lost heat. When we feel the beaker as hot/cold we are feeling the surroundings. Used for phase change NO temp. change

Reactions that Release Energy These reactions are Exothermic - q [negative value] the system loses heat The surroundings heat up

Reactions that Absorb Energy These reactions are Endothermic + q [positive value] system gains heat the surroundings cool down

Graphing Heat in Reactions We’ve already looked at one graphical representation… q = mHv q = mCΔT Temperature q = mHf Heat (Joules) or Time

What Heating Curves Tell Us Heating Curves tell us how temperature changes when energy is absorbed or released. For Example: Energy Added  Kinetic Energy  ΔTemp Energy Added  Potential Energy  Δ Phase Heating Curves do NOT tell us how much energy was needed to get the rxn started. Heating Curves do NOT tell us how the energy used to start the reaction compares to the energy produced by the reaction.

However, only a small fraction of collisions produces a reaction. Why? We know that molecules must collide to react. However, only a small fraction of collisions produces a reaction. Why? Particles must collide with the correct orientation and the right amount of energy Particles lacking the correct site or necessary kinetic energy to react bounce apart unchanged when they collide.

How Reactions Happen Activation energy is the amount of energy needed to cause a reaction. Collisions must equal or exceed the activation energy Reactions can occur: Very fast – such as a firecracker Very slow – such as the time it took for dead plants to make coal In chemistry, reaction rate is measured in terms of moles of reactant consumed (or product formed) per liter per second

Potential Energy Diagrams A complete picture of the system’s net change in energy (HEAT of REACTION) is best illustrated by a POTENTAL ENERGY DIAGRAM In a Potential Energy Diagram potential energy of the reactants and products is drawn with reference to a time or rate sequence. ΔH = q, (HEAT of Reaction (AKA enthalpy) ΔH = Energy Contained in of Products – Energy contained in of Reactants

Ways to Express ΔH – Heats of Reaction Endothermic Reactions ΔH is a positive value Example – N2 + O2 2NO ΔH = 182.6 kJ ΔH can be written as a reactant in an equation Example - N2 + O2 + 182.6 kJ 2NO Exothermic Reactions 1. ΔH is negative value Example – 2CO + O2 2CO2 ΔH = - 566.0 kJ 2. ΔH can be written as a product in an equation Example – 2CO + O2 2CO2 + 566.0 kJ

Practice Problem

http://player. discoveryeducation. com/index. cfm http://player.discoveryeducation.com/index.cfm?guidAssetId=AFD8D965-FC9F-4B48-89BB-B249027F3E7C&blnFromSearch=1&productcode=US Standard Deviants Chapter 17 – Chemistry Videos - Reaction Energy Diagrams (2:27min) Chapter 17 – Chemistry Videos – Exothermic & Endothermic (2:39min), Chapter 17 – Chemistry Videos - Transition Time & Energy (1:23 min)

Table I: Heats of Reaction kJ = 1000 Joules Exo- thermic Endo- thermic DH = ( PE of Products) - (PE of Reactants)

2 2

A POTENTIAL ENERGY DIAGRAM: C + O2 → CO2 + 395 KJ Products have lower energy than reactants Energy released Exothermic ΔH = -395kJ (Heat of Reaction) C + O2 Energy 395kJ CO2 Reactants ® Products

A POTENTIAL ENERGY DIAGRAM: CaCO3 + 176kJ → CaO + CO2 Products have higher energy than reactants Energy absorbed Endothermic ΔH = +176kJ (Heat of Reaction) CaO + CO2 Energy 176kJ CaCO3 Reactants ® Products

Potential Energy Diagrams One thing that is not represented in these simple examples… Activation energy Minimum amount of energy required to start a reaction

Endothermic with Activation Energy However, adding 150kJ won’t make the reaction happen 100 kJ 250 kJ Products have 350kJ of PE Activation Energy ΔH = Heat of Reaction 150 kJ Reactants have 200kJ of PE

Endothermic with Activation Energy Activated Complex – temporary, unstable, intermediate union of reactants 100 kJ 250 kJ 150 kJ 450 kJ ΔH = 150kJ Activation Energy

Exothermic with Activation Energy Activation energy = 10kJ Activation energy Reactants 40kJ ΔH = 25kJ ΔH Products15kJ

Catalyst: something that reduces the amount of activation energy needed to start a reaction – speeds up a reaction Activation energy with a catalyst

Endothermic Reaction with a Catalyst

Exothermic Reaction with a Catalyst

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“Reaction Rates and Equilibrium”

Five Factors Affecting Reaction Rate Nature of the Reactants: what the reactants are (solids, liquids or gases) Temperature: Increasing temperature always increases the rate of a reaction. Surface Area: Increasing surface area increases the rate of a reaction

Five Factors Affecting Reaction Rate Concentration: Increasing concentration USUALLY increases the rate of a reaction Presence of Catalysts Catalyst: A substance that speeds up a reaction, without being consumed itself in the reaction

Reversible Reactions SOME reactions do not go to completion These reactions may be reversible a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously Forward: 2SO2(g) + O2(g) → 2SO3(g) Reverse: 2SO3(g)  2SO2(g) + O2(g)

Reversible Reactions 2SO2(g) + O2(g) ↔ 2SO3(g) The two equations can be combined into one, by using a double arrow, which tells us that it is a reversible reaction: 2SO2(g) + O2(g) ↔ 2SO3(g)

Reversible Reactions For a reversible reaction, the rates of the forward and reverse reactions are equal Thus a chemical equilibrium occurs A dynamic condition in which the forward and reverse reactions are proceeding at equal rates, producing an apparent constant condition or no net change in the actual amounts of the components of the system Equilibrium is recognized when observable changes such as color, temperature or pressure no longer occur

2SO2(g) + O2(g) ↔ 2SO3(g) What actually happens when sulfur dioxide and oxygen gas are mixed in a sealed container? Initially the two reactants begin to react to form the product but no sulfur trioxide exists at the beginning of the reaction The initial rate of the reverse reaction is zero As product is formed the decomposition of sulfur trioxide begins The reverse reaction proceeds slowly at first but it’s rate increases as the concentration of sulfur trioxide increases Simultaneously the rate of the forward reaction decreases because sulfur dioxide and oxygen are being used up Eventually, the synthesis and decomposition occur at the same rate Equilibrium is reached Both forward and reverse reactions continue at the same rate and the concentrations of the reactants and products remains the same or CONSTANT

Reversible Reactions Equilibrium is reached Both forward and reverse reactions continue at the same rate and the concentrations of the reactants and products remains the same This does not necessarily mean the concentrations of the components on both sides of the chemical equation are equal In fact the concentrations can be dramatically different It means the concentrations of the components remain constant

Factors Affecting Equilibrium Chemical equilibrium is a delicate balance Changes of almost any kind disrupt the balance When the equilibrium of a system is disrupted the system makes adjustments to restore the equilibrium However the equilibrium position of the restored equilibrium is different than the original equilibrium position The amounts of reactants or products may have increased or decreased The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shifts as a result of changing conditions

Le Chatelier’s Principle If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress Stresses on the equilibrium are considered to be: Changes in concentration of reactants or products Changes in temperature Changes in pressure (only if a gas is present)

Le Chatelier’s Principle Concentration Increasing the concentration of one substance in a reaction at equilibrium will cause the reaction to proceed in the direction that will use up the increase Adding more reactant produces more product Removing the product as it forms will produce more product

Le Chatelier’s Principle Example – Concentration Change Decomposition of carbonic acid in aqueous solution to form the products carbon dioxide and water Add CO2 Direction of shift H2CO3(aq) CO2(aq) + H2O(l) Direction of shift Remove CO2

Le Chatelier’s Principle Temperature Increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat – endothermic Decreasing the temperature causes the equilibrium to shift in the direction that releases heat – exothermic If heat is one of the products – exothermic reaction Cooling an exothermic reaction will produce more product Heating an exothermic reaction will shift the reaction to the reactant side of the equilibrium

Le Chatelier’s Principle Example – Temperature Change Sulfur trioxide is produced in an exothermic reaction of sulfur dioxide and oxygen Add heat Direction of shift 2SO2(g) + O2(g) 2SO3 + heat Direction of shift Remove heat (cool)

Le Chatelier’s Principle Pressure Changes in pressure will only effect gaseous equilibria Increasing the pressure will usually favor the direction that results in a smaller volume caused by fewer molecules

Le Chatelier’s Principle Example – Pressure Change N2(g) + 3H2(g) 2NH3(g) Increase pressure Direction of shift Direction of shift Decrease pressure For every two molecules of ammonia made, four molecules of reactant are used up – the equilibrium shifts to the right with an increase in pressure and shifts to the left with a decrease in pressure

N H Equilibria in the Real World –The Haber Process – Production of Ammonia N H Major Uses Fertilizers (for example ammonium nitrate (NH4NO3) Synthesis of Nitric Acid (HNO3) Dyes Drugs Structure is 1 Nitrogen atom to 3 Hydrogens. Oxidation (combustion, burning of ammonia makes nitric acid) Ammonium nitrate is made by neutralisation of nitric acid with ammonia.

Haber Process - Gas Phase Equilibrium catalyst N2(g) + 3H2(g) <=====> 2NH3(g) + heat high pressure and temperature ΔH= - 92 kJ See Reference Table I

The Principle of Le Chatelier Changes in Concentration or Pressure N2(g) + 3 H2(g) 2 NH3(g) an increase in N2 and/or H2 concentration or pressure, will cause the equilibrium to shift towards the production of NH3

The Principle of Le Chatelier Changes in Concentration or Pressure N2(g) + 3 H2(g) 2 NH3(g) likewise, a decrease in NH3 concentration will cause more NH3 to be produced A decrease in pressure will cause more N2 and H2 to be produced

The Principle of Le Chatelier Changes in Temperature N2(g) + 3 H2(g) 2 NH3(g) + heat for an exothermic reaction, an increase in temperature will cause the reaction to shift back towards reactants Increase in temperature favors the endothermic reaction

Practice Problem The cobalt complexes participating in the equilibrium below comprise a humidity sensor. From Le Châtelier's principle, when the sensor is moist (excess H2O), what color is the cobalt complex? pink, blue

Hb(O2)4 + 4 CO (g)   Hb(CO)4 + 4O2 (g) Practice Problem A competition experiment involves O2 and CO vying for hemoglobin (Hb) sites, defined by the equilibrium Hb(O2)4 + 4 CO (g)   Hb(CO)4 + 4O2 (g) From Le Châtelier's principle, how is CO poisoning reversed? decrease O2 pressure, increase O2 pressure, remove Hb(CO)4

Energy in Reactions Things in Nature move toward: greater stability thus lower energy (exothermic energy changes) less organization MORE CHAOS (increased randomness)

1. Energy Changes Systems in condition of great energy tend to change to conditions of low energy. Thus exothermic reactions are favored in nature

2. Randomness or Entropy (S) Entropy is a measure of disorder or randomness, Represented by the symbol S Systems tend to change from conditions from of great order to conditions of low order. Thus entropy changes from low to high. Your room NEVER cleans itself (disorder to order?)

Entropy When is there more disorder in a system? 1. Entropy of gases is greater than the entropy of liquids and solids entropy increases in reactions in which solid reactants form liquids or gaseous products S l g ΔS

Entropy 2. Entropy increases when a substance is divided into parts Entropy increases when ionic compounds dissolve in water Decomposition reactions 3. Entropy increases in chemical reactions in which total number of product molecules is greater than the total number of reactant molecules

Entropy 4. Entropy tends to increase when temperature increases As temperature increases molecules move faster and faster which increases disorder

LeChatlier Principle Animation: http://www. mhhe