Chemistry SAC 2: Rates and Equilibrium

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Chem SAC 2 Flashcards

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48 Terms

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Rate of reaction

The change in concentration of reactants/products over time, measured in mol/L/s

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Collision theory

For a reaction to occur, reactant particles must:

  • Collide with each other

  • Have sufficient energy to break reactant bonds

  • Collide in the correct orientation to break reactant bonds

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Activation energy (Ea)

The minimum energy required to break reactant bonds in a collision.

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High vs. low activation energy

  • High: strong reactant bonds, lower proportion of particles have the minimum energy required to react.

  • Low: weak reactant bonds, higher proportion of particles have the minimum energy required to react.

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Transition state

A new arrangement of atoms when activation energy is absorbed - unstable because reactants break and products form

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Reaction rate graphs

Graphs which measure amount of product/reactant vs. time

<p>Graphs which measure amount of <strong>product/reactant</strong> vs. time</p>
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Reaction rate graph terms

  • Gradient: the rate of reaction

  • Tangent: the rate of reaction at any instant

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Reaction rate conditions

  • Surface area of a solid reactant

  • Concentration of a dissolved reactant

  • Pressure of a gaseous reactant

  • Temperature

  • Presence of a catalyst

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Concentration vs. reaction rate

  • Concentration: more dissolved reactant particles per unit volume

  • Increasing conc —> increasing frequency of collisions —> increasing successful collisions with correct orientation

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Pressure vs. reaction rate

  • Pressure: adding more reactant gas/decreasing container volume

  • Increasing press —> increasing frequency of collisions —> increasing successful collisions with correct orientation

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Surface area vs. reaction rate

  • Surface area: more particles are exposed at the surface when solids are broken down.

  • Increasing SA —> increasing frequency of collisions —> increasing successful collisions with correct orientation

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Temperature vs. reaction rate

  • Temperature increases energy of collisions —> increased collisions with >= activation energy —> increases successful collisions

  • Temperature causes faster-moving particles —> increases frequency of collisions —> increases successful collisions

  • Energy increase has a greater effect than frequency of collisions on reaction rate

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Maxwell-Boltzman distribution

  • A graph that shows the kinetic energy range of particles in a substance

  • Higher temperatures lead to an increase in an energy, hence a much greater proportion of particles have >= activation energy

<ul><li><p>A graph that shows the kinetic energy range of particles in a substance</p></li></ul><ul><li><p>Higher temperatures lead to an increase in an energy, hence a <strong>much greater proportion</strong> of particles have &gt;= activation energy</p></li></ul><p></p>
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Homogenous catalyst

Catalysts in the same physical state as the reactants & products

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Heterogenous catalyst

Catalysts in a different physical state as the reactants & products

  • More easily separated from products

  • Reused easily

  • Often used at high temps

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Catalysts vs. reaction rate

  • Greater proportion of reactant collisions >= activation energy —> leads to a chemical change

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Irreversible chemical reaction

A reaction that proceeds in one direction, indicated by one-way arrows

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Reversible chemical reaction

A reaction that goes in two directions, where products can reform reactants.

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Open system

A system where matter and heat energy can be exchanged with the surroundings.

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Closed system

A system where matter cannot be exchanged with the surroundings, but heat energy can.

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Equilibrium

The rates of both forward and reverse reactions are equal

  • Concentrations and amounts are constant

  • Temperature is constant

  • Mixture of both reactants and products

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Dynamic equilibrium

Both forward and reverse reactions are occurring at the same rate —> bonds constantly break/reform

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Homogenous equilibria

An equilibrium system that has all reactants and products in the same state

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Heterogenous equilibria

An equilibrium system that has reactants and products in different states

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Rate-time graph

A graph where equilibrium is reached and when the forward & reverse reaction rate is the same.

<p>A graph where <strong>equilibrium is reached</strong> and when the forward &amp; reverse reaction rate is the same.</p>
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Concentration-time graph

A graph where equilibrium is reached and when the concentration of all species is constant

<p>A graph where <strong>equilibrium is reached</strong> and when the concentration of all species is <strong>constant</strong></p>
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Extent of reaction

The relative amount of reactants and products present at a particular time, and does not give information on how fast the reaction proceeds.

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Reaction quotient (Q)

  • A numerical measure of the extent of a reaction

  • For the equation aW + bX <=> cY + dZ, Q is calculated by:

    • ([Y]^c * [Z]^d) / ([W]^a + [X]^b)

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Equilibrium constant (K)

  • The reaction quotient (Q) at equilibrium (fixed temp)

  • For the equation aW + bX <=> cY + dZ, K is calculated by:

    • ([Y]^c * [Z]^d) / ([W]^a + [X]^b)

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Units for equilibrium constants & reaction quotients

Use M (mol/L) for units:

  • For the example [NH3]² / [H2]³ * [N2], the units should be M^-2 (from 2 - (3 + 1) through index laws)

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Q & K: relative reactants and products

  • Q/K < 10^-4: value of Q/K is small which suggests that there are mostly reactants in a mixture

  • 10^-4 < Q/K < 10^4: value of Q/K is moderate which suggests that there are a significant amount of both reactants and products.

  • Q/K > 10^4: value of Q/K is large which suggests that there are mostly products in a mixture

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Q vs. K

  • Q > K: the reaction shifts to the ‘left’ and favours the reverse reaction to reach equilibrium (by forming more reactants)

  • Q < K: the reaction shifts to the ‘right’ and favours the forward reaction to reach equilibrium (by forming more products)

  • Q = K: the reaction is at equilibrium and does not shift.

<p></p><ul><li><p><strong>Q &gt; K</strong>: the reaction shifts to the <strong>‘left’</strong> and favours the <strong>reverse reaction</strong> to reach equilibrium (by forming <strong>more reactants</strong>)</p></li><li><p><strong>Q &lt; K</strong>: the reaction shifts to the <strong>‘right’</strong> and favours the <strong>forward reaction</strong> to reach equilibrium (by forming <strong>more products</strong>)</p></li><li><p><strong>Q = K</strong>: the reaction is at <strong>equilibrium</strong> and does not shift.</p></li></ul><p></p>
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K in different equations

  • if one reaction is the reverse of another, the two equilibrium constant (K) values are reciprocal of each other

  • If the coefficients of a reaction are doubled, the value of K is squared.

  • If the coefficients of a reaction are halved, the value of K is square rooted.

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RICEC tables

  • R = Reaction, I = Initial, C = Change, E = Equilibrium, C = Concentration

    • 1. R: Write the reaction down

    • 2. I: Identify the initial mole amounts of reactants and products (0)

    • 3. C: Calculate the change in the reactants (-) and products (+) according to the mole ratio

    • 4. E: Minus the I and C row to calculate the amounts at equilibrium

    • 5. C: Calculate the concentrations using n/V and hence, find K (equilibrium constant)

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Position of equilibrium

The relative amounts of reactants and products at equilibrium. It is affected by:

  • Adding/removing reactants/products

  • Changing pressure by changing volume

  • Dilution (aqueous equilibria)

  • Changing temperature

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Le Chatelier’s principle (LCP)

If a system at equilibrium is subjected to a change, a net reaction will occur that partially opposes the change and the system will establish a new equilibrium.

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Adding reactants/removing products

  • The reaction shifts to the ‘right’ and favours the forward reaction to consume reactants/increase products

  • Eg. N2 + 3H2 —> 2NH3: the new concentration of N2 is still higher than its original concentration, the change was partially opposed.

  • K does not change with concentration

<ul><li><p>The reaction shifts to the <strong>‘right’</strong> and favours the <strong>forward reaction</strong> to consume reactants/increase products</p></li><li><p>Eg. N2 + 3H2 —&gt; 2NH3: the new concentration of N2 is still higher than its original concentration, the change was <strong>partially opposed</strong>.</p></li><li><p>K does not change with <strong>concentration </strong></p></li></ul><p></p>
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Adding products/removing reactants

  • The reaction shifts to the ‘left’ and favours the reverse reaction to partially oppose the change (decrease products/increase reactants)

  • Shifts in concentration follow the mole ratio in the balanced equation.

  • K does not change with concentration

<ul><li><p>The reaction shifts to the <strong>‘left’</strong> and favours the <strong>reverse reaction</strong> to partially oppose the change (decrease products/increase reactants)</p></li><li><p>Shifts in <strong>concentration</strong> follow the <strong>mole ratio</strong> in the balanced equation.</p></li><li><p>K does not change with <strong>concentration </strong></p></li></ul><p></p>
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Increasing volume/decreasing pressure

The reaction shifts to the side with more particles and concentrations drop at the new equilibrium

<p>The reaction shifts to the side with <strong>more particles</strong> and <strong>concentrations drop</strong> at the new equilibrium</p>
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Decreasing volume/increasing pressure

The reaction shifts to the side with less particles and concentrations remain higher at the new equilibrium

<p>The reaction shifts to the side with <strong>less particles</strong> and <strong>concentrations</strong> <strong>remain higher</strong> at the new equilibrium</p>
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Volume changes using equilibrium law

If volume is halved, then concentration will double (c = n/v), leading to K > Q, hence a net forward reaction must occur to partially oppose the change.

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Changing pressure by adding inert gases

  • The addition of an inert gas (eg. He) will increase overall pressure in a gaseous equilibrium

  • The concentrations of the reactants remain the same, so there is no change in equilibrium position

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Dilution (adding water)

  • Adding water reduces the number of dissolved particles per unit volume

  • Hence the reaction shifts to the side with more dissolved particles

  • Eg. Fe3+ + SCN- —> FeSCN2+ dilution would result in a reverse reaction as the reactants have 2 particles compared to 1.

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Temperature increase vs. K

  • Temperature increase means that heat energy is added to a reaction, and the system must consume heat energy to partially oppose the change.

    • Endothermic: shifts right, more products, K increases

    • Exothermic: shifts left, more reactants, K decreases

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Temperature decrease vs. K

  • Temperature decrease means the system moves in the direction that releases heat energy

    • Endothermic: shifts left, more reactants, K decreases

    • Exothermic: shifts right, more products, K increases

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Example of changing temperature

Eg. N2 + O2 —> 2NO deltaH = +180kJ

  • A temp increase will shift the reaction to the right and K increases

<p>Eg. N2 + O2 —&gt; 2NO deltaH = +180kJ</p><ul><li><p>A temp increase will shift the reaction to the <strong>right </strong>and K increases</p></li></ul><p></p>
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Catalysts

Catalysts provide an alternate reaction pathway by lowering activation energy, hence a greater proportion of particles have the energy needed for successful collisions.

  • Increases both forward & reverse reactions

  • Do not change equilibrium position or K

  • Increase the rate at which equilibrium is established.

<p>Catalysts provide an <strong>alternate reaction pathway</strong> by <strong>lowering activation energy</strong>, hence a greater proportion of particles have the energy needed for successful collisions.</p><ul><li><p>Increases both forward &amp; reverse reactions</p></li><li><p>Do not change equilibrium position or K</p></li><li><p>Increase the rate at which equilibrium is established.</p></li></ul><p></p>
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Yield

The amount of desired product from a chemical reaction.