CHEMICAL EQUILIBRIUM EQUILIBRIUM CONSTANT LE ... · 7. 4. 2020 1 •• acids & bases ••...

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•• ACIDS & BASESACIDS & BASES•• ACIDACID--BASE EQUILIBRIUMBASE EQUILIBRIUM•• BUFFERSBUFFERS

•• THERMODYNAMICSTHERMODYNAMICS

•• CHEMICAL EQUILIBRIUMCHEMICAL EQUILIBRIUM•• EQUILIBRIUM CONSTANTEQUILIBRIUM CONSTANT•• LE CHLE CHÂÂTELIERTELIER´́S PRINCIPLES PRINCIPLE

06-03-2020

(Zuzana Országhová)

CHEMICAL EQUILIBRIUMCHEMICAL EQUILIBRIUM

EQUILIBRIUM CONSTANTEQUILIBRIUM CONSTANT

LE CHLE CHÂÂTELIERTELIER´́S PRINCIPLES PRINCIPLE

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The Concept of Equilibrium

N2O4 (g) 2 NO2 (g)forward

reverse

reactant product

reactantproduct

• At equilibrium, the forward and reverse reactions are proceeding at the same rate.

• Once equilibrium is achieved, the amount(concentration) of each reactant and product remains constant.

The Concept of Equilibrium

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• Consider the reaction:

• The equilibrium expression for this reaction would be

Keq = [C]c[D]d

[A]a[B]b

aA + bB cC + dD

The Equilibrium Constant

Reactants Products�

K << 1

Reactants Products�

K >> 1

What What ddoesoes the the vvaluealue of of KK mmeanean??

• If K >> 1, the reaction is product-favored; product predominates at equilibrium.

(„eq. lies to the right“)

• If K << 1, the reaction is reactant-favored; reactant predominates at equilibrium.

(„eq. lies to the left“)

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The equilibrium between reactants and products may be disturbed in three ways:

(1) by changing the concentration of a reactant or product

(2) by changing the temperature(3) by changing the volume (for systems involving gases)

A change in any of these factors will cause a system at equilibrium to response - Le Le ChatelierChatelier’s’s principleprinciple..

Disturbing a Chemical EquilibriumDisturbing a Chemical Equilibrium

Le Châtelier’s PrincipleLe Châtelier’s Principle

“If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.”

•The Le Châtelier's principle states that:if a system at equilibrium is disturbed, the position of equilibrium will shift to minimize the disturbance.

Henry Louis Le Chatelier(1850-1936)

„All my life I maintained respect for order and law. Order is one of the most perfect forms of civilization“

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What do you do when your stomache growls from hunger?• Feed it!

What do you do when your checking account is at a zero balance?• Ask your mom to deposit more money, of course!

• When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away.

• When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

Simply said.....Simply said.....

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The Effect of Changes in The Effect of Changes in ConcentrationConcentration

Consider general reaction:

Change Shifts the Equilibrium

Increase concentration of product(s) left

aA + bB cC + dD

Add

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Decrease concentration of product(s) right

aA + bB cC + dD

Remove





Increase concentration of reactant(s) right

aA + bB cC + dD

Add

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Decrease concentration of reactant(s)

Increase concentration of reactant(s) right

left

aA + bB cC + dD

Remove

The Le Châtelier principle in physiologyThe Le Châtelier principle in physiology

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OOxygenxygenTransport Transport

Hb + O2 � HbO2

• The equilibrium shifts to the right

• Hb and O2 combine to make more HbO2

OO22 Transport Transport

• The equilibrium shifts to the left

• HbO2 breaks down (dissociates) increasing the amount of free O2.

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The Effect The Effect of of TemperatureTemperature on Equilibriumon Equilibrium

• Exothermic reactions release energy (heat) and endothermic reactions absorb energy(heat).

• Writing heat as a product in an exothermic reaction or as a reactant in an endothermic reaction - helps us use Le Châtelier’s principle to predict the effect of temperature changes

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The Effect of The Effect of TemperatureTemperature Changes Changes ononEquilibrium for Equilibrium for Exothermic Exothermic ReactionsReactions

• For an exothermic reaction - heat is a product.• Increasing the temperature is like adding product.• Decreasing the temperature is like removing product.



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The Effect of The Effect of TemperatureTemperature Changes Changes ononEquilibrium for Equilibrium for Endothermic Endothermic ReactionsReactions

• For an endothermic reaction, heat is a reactant• Increasing the temperature is like adding reactant.• Decreasing the temperature is like removing reactant.

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−196°C 0°C 23°C 35°C 50°C

Equilibrium Exercise #Equilibrium Exercise #22

Determine whether the following reactions are favoredby high or low temperatures?

1. 2SO2(g) + O2(g) ⇄ 2 SO3(g); ∆Ho = -180 kJ/mol

2. CO(g) + H2O(g) ⇄ CO2(g) + H2(g); ∆Ho = -46 kJ/mol

3. CO(g) + Cl2(g) ⇄ COCl2(g); ∆Ho = -108 kJ/mol

4. N2O4(g) ⇄ 2 NO2(g); ∆Ho = +57 kJ/mol

5. CO(g) + 2H2(g) ⇄ CH3OH(g); ∆Ho = -270 kJ/mol

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CCheckinghecking comprehensioncomprehension

Consider the following equilibrium.

H2O2(l ) � H2(g) + O2(g) ∆H = 187 kJ/mol

Which way will the equlibrium shift if:

1) More H2 is added.left (use up H2)

2) O2 is removed.right (restore O2)

3) The temperature increases.right (endothermic, heat is a reactant)

ACIDS & BASESACIDS & BASES

ACIDACID--BASE EQUILIBRIUMBASE EQUILIBRIUM

BUFFERSBUFFERS

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Properties of BASESProperties of BASES• Have a slippery feeling• Taste Bitter• Corrosive• Can conduct electricity• Do not react with metals• Turns red litmus paper blue

Properties of ACIDSProperties of ACIDS• Have a sour taste• Corrosive• Conduct electricity• Some acids react strongly with metals• Turns blue litmus paper red

Arrhenius acids and basesArrhenius acids and bases

The fundamental concept:

Acid - any substance which delivers hydrogen ion (H+) to the solution.

HA → H+ + A¯

Base - any substance which delivers hydroxide ion (OH¯) to the solution.

BOH → X+ + OH¯

Svante August Arrhenius proposed that substances exists as ions in solution in his dissertation, which was awarded a fourth class (D) in 1884. He was unable to find a job in his native Sweden. He was awarded the Nobel Prize in 1903 for his electrolytic dissociation theory.

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BrønstedBrønsted--Lowry Lowry theory theory of of acids and basesacids and bases

An acid is a substance from which a proton can be removed.Acids are proton donors.

A base is a substance that can remove a proton from an acid.Bases are proton acceptors.

Acid-base reactions are competitions for protons.HCl + H2O ↔ H3O

+ + Cl-

HNO3 + H2O ↔ H3O+ + NO3

-

CH3COOH + H2O ↔ H3O+ + CH3COO-

H2O + NH3 ↔ NH4+ + OH-

conjugate acids and bases

SelfSelf--ionization of waterionization of water

Water molecules autoionize:

H2O + H2O ←→ H3O+ + OH¯

[H3O+]×[OH¯]

Keq = ————————[H2O]2 ([H2O] = 1000/18 = 55.6)

Kw = [H2O]2 × Keq = [H3O+]×[OH¯]= 1×10-14 [mol2.l-2]

(only at 25oC, it’s T dependent)

The ion product of water, Kw increases as T increases, and its value remains the same in the presence of acid or base.

3.24 .×10-18

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Strong Strong acids acids vs. Weak acidsvs. Weak acids

100% dissociation into ions

partial dissociation, both ions and molecules

The pH scaleThe pH scale

Søren Peder Lauritz Sørensen introduced the pH scale in 1909. The p is from the German word potenz,power of (10).

pH = – log [H+]; [H+] = 10 –pH

pOH = – log [OH–]; [OH–] = 10 –pOH

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pH = − log [H3O+]

pOH = − log [OH-]

pH + pOH = 14 pH + pOH = 14

3 situations can exist for the relative concentrations of hydronium and hydroxide ions in water:•in neutral solution, [H3O

+] = [OH-] = 10-7 mol.l¯1

•in acidic solutions, [H3O+] > [OH-] pH < 7

•in basic solutions, [H3O+] < [OH-] pH > 7

acidity constant

HA + H2O ←→ H3O+ + A-

↑ ↑

Weak acid Conjugate base

Ionization of weak acidsIonization of weak acids

The pKa is defined similar to the pH: pKa = – log Ka

Ka = 10 –pKa

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Calculation of pHCalculation of pH –– formulas:formulas:

strongstrong acidacid pH = − log [ H3O+]

strongstrong basebase pOH = − log [OH-]pH = 14 − pOH

weakweak acidacid pH = ½ pKa − ½ log ca

weakweak basebase pOH = ½ pKb − ½ log cb

pH = 14 − ½ pKb + ½ log cb

[OH-]

[H+] pOH

pH

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pHpH

1 2 3 4 5 6 8 9 10 11

The biological view in the human bodyThe biological view in the human body

acidic basic/alkaline7

Tortora & Grabowski, Prin. of Anatomy & Physiology, 10th ed., Wiley (2003)

Small changes in pH can produce Small changes in pH can produce major disturbancesmajor disturbances

• Change of enzymes activity

• Change of metabolism

• Effect on electrolytes

• Failure of function of excitable tissues (muscle, heart, nervous system)

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Control of Control of pH in organismpH in organism

• chemical buffer systems (buffers)

• respiratory system (lungs)

• urinary system (kidneys)

BUFFER SYSTEMSBUFFER SYSTEMS

Composition:

• Mixture of a weak acid and salt of its conjugate base

• Mixture of a weak base and salt of its conjugate acid

Function:

• maintaining the pH of a solution constant by taking up protons that are released during reactions, or by releasing protons when they are consumed by reactions

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Composition:



Function:

• maintaining approximately constant pH upon small additions of acid or base.


Composition:



Function:

• maintaining approximately constant pH upon small additions of acid or base.

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Mechanism of Mechanism of buffer buffer actionaction

Addition of a strong acid (HCl):

CH3COOH + CH3COO- + Na+ +HCl 2CH3COOH + NaCl

Addition of a strong base (NaOH):

CH3COOH + CH3COO- + Na+ + NaOH 2CH3COO- + 2Na+ + H2O

ACETATE BUFFER = Mixture of a weak acid and salt of its

conjugate base = CH3COOH and CH3COO-Na+

Buffer Buffer capacitycapacity

= the quantity of strong acid (or base) that has to be added to one liter of a buffer solution to change its pH value by one unit.

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CalculationCalculation of buffer pHof buffer pH(Weak acid + salt of its conjugate base)

HA + H2O ←→ H3O+ + A-

↑ ↑

Weak acid Conjugate base

CalculationCalculation of buffer pHof buffer pH(Weak acid + salt of its conjugate base)

Henderson-Hasselbalch equation

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BuffersBuffers in in biologicalbiological systemssystemsand and theirtheir primaryprimary rolesroles

Buffer system Major function

Bicarbonate(H(H22COCO33 / HCO/ HCO33

--))Primary ECF buffer against non-carbonic acid changes

Protein((protein / proteinateprotein / proteinate))

Primary ICF buffer; also buffers ECF

Hemoglobin((Oxyhemoglobin / Oxyhemoglobin / hemoglobinhemoglobin))

Primary buffer against carbonic acidchanges

Phosphate((HH22POPO44

-- / HPO/ HPO4422--))

Important urinary buffer; also buffersICF

EffectEffect of of bicarbonatebicarbonate buffer buffer systemsystem

H2O + CO2 H2CO3 HCO3– + H+

(carbonic anhydrase)

NaHCO3 / H2CO3(salt) (acid)

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H+ Cl- + NaHCO3 NaCl + H2CO3

+ H2CO3 + H2CO3

Na+ OH- + H2CO3 NaHCO3 + H2O

+ NaHCO3 + NaHCO3


(salt)

(acid)


H+ Cl- + NaHCO3 NaCl + 2 H2CO3

+ H2CO3

Na+ OH- + H2CO3 2 NaHCO3 + H2O

+ NaHCO3


(salt)

(acid)

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• pKa = 6.1• [HCO3

-] = 24 mmol.l-1

• [H2CO3] = 1.2 mmol.l-1

[HCO3-]

[H2CO3]= 20

[NaHCO3]pH = pKa + log

[H2CO3]

Henderson – Hasselbalch equation

BicarbonateBicarbonate buffer buffer systemsystem

[NaHCO3]pH = 6.1 + log

[H2CO3]


lungs

kidneys

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[NaHCO3]pH = 6.1 + log

[CO2]aq


lungs

kidneys

Bicarbonate buffer systemBicarbonate buffer system

246.1 + log = 7.4

1.2

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2 HCO3- + 2H+ 2CO2 + 2H2O

226.93 = 6.1 + log

3.2

227.36 = 6.1 + log

1.2

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HemoglobinHemoglobin buffer buffer systemsystem

acidic metabolites

HHb H+ + Hb- pK = 7.82

in lungs +O2 in tissue -O2

HHbO2 H+ + HbO2- pK = 6.17

+ HCO3- H2CO3 H2O

CO2

exhaled

Oxyhemoglobin is stronger acid than deoxyhemoglobin.

ThermodynamicsThermodynamics

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ThermodynamicsThermodynamics(Greek: thérme-, “heat”; dy’namis, “power”)• the branch of physical science that deals with the relations between heat and other forms of energy (such as mechanical, electrical, or chemical energy)

What is energy?

„... the term energy s difficult to define precisely, but one possible definition might be the capacity to produce an effect“

Encyclopaedia Britannica

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What is energy?

„... the term energy s difficult to define precisely, but one possible definition might be the capacity to produce an effect“

Encyclopaedia Britannica

Energy is commonly defined as the capacity to do work or transfer heat.

Work is the energy used to cause an object to move against a force, and heat is the energy used to cause the temperature of an object to increase.

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• Heat and work are modes of transfer of energy and not ‘energy’ itself.

Once inside the system, the part which came via work and the part which came via heat, cannot be distinguished*.

• From the above it is clear that, bodies contain internal energy and not heat (nor work!).

� All objects (also molecules) can possess kinetic energy, the energy of motion.

� The magnitude of the kinetic energy, Ek, of an object depends on its mass, m, and speed, v:

Ek = ½ mv2

� All other kinds of energy - the energy stored in a stretched spring, or e.g. in a chemical bond - are potential energy. An object has potential energyby virtue of its position relative to other objects.

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The potential energyinitially stored in the motionless bicycle at the top of the hill is converted to kinetic energy as the bicycle moves down the hill and losespotential energy.

Work and energy consist of two quantities: an intensity factor (driving force), and capacity factor (measure of substance quantity).

ReactantsCurrent

sourceElevated position

The forms of work

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The „language“ of thermodynamics (TD)

• System is region where we focus our attention

• Surrounding is the rest of the universe

• Universe = System + Surrounding

system

surroundings

•no exchange of matter•no exchange of heat

heat

system

surroundings

•no exchange of matter•can exchange heat energy

system

surroundings

matter

heat

•can exchange matter•can exchange heat energy

Types of termodynamic systems

isolated closed open

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State functions in TD

• A property which depends only on the current state of the system (as defined by T, P, V etc.) (initial and final) and not on the way, nature and pathways by which the system gets into the final state.

• A process for which the final and initial states are the same is called a cyclic process. For a cyclic process change in a state function is zero.

Spontaneous and Driven processes

• A spontaneous process is one which occurs

‘naturally’, ‘down-hill’ in energy. (The process does

not require input of work in any form to take place.)

Melting of ice at 50°C is a spontaneous process.

• A driven process is one which wherein an external

agent takes the system uphill in energy (usually by

doing work on the system).

Freezing of water at 50°C is a driven process (you need a

refrigerator, wherein electric current does work on the

system).

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Laws of thermodynamics

can help to predict the feasibility of a physical

process, its direction, to understand why

molecules adopt their natural conformation or why

move through cell membranes.

But thermodynamics says nothing about the speed of processes!

The first law of thermodynamics

• In any physical or chemical change, the total energy of a system, including its surroundings, remains constant (principle of conservation of energy).

• Energy is never created nor destroyed....and so

...any energy that is lost by a system must begained by the surroundings, and vice versa.

ΔU = Q + W

Q - the heat exchanged between a system and its surroundings (added to or liberated from the system)

W - the work done on or by the system

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Sign Conventions for Q, W and ∆U

Q+ system gains

heat- system loses

heat

W+ work done

on system- work done

by system

∆U+ net gain of

energy by system

- net loss of energy by system

� Definition of internal energy, U, of a system = the sum of all the kinetic and potential energies of the components of the system.

For the system example, the internal energy includes not only the motions and interactions of the H2 and O2 molecules but also the motions and interactions of the nuclei and electrons.

� We generally do not know the numerical value of a system’s internal energy.

� In thermodynamics,we are mainly concerned with the change in U (in other quantities as well).

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cooling warming

The internal energy of the waterat 25 °C is the same in either case

The internal energy The internal energy -- a state functiona state function

Hess’s lawthe heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps

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� A process in which the system absorbs heat, is called endothermic (endo- means “into”).

During an endothermic process, such as the melting of ice, heat flows into the system from its surroundings

� A process in which the system loses heat is called exothermic (exo- means “out of ”).

During an exothermic process, such as the combustion of gasoline, heat exits or flows out of the system into the surroundings.

Enthalpy (H) = the heat content of a system

The mathematical formulation of enthalpy:

H = U + pV

H - enthalpy of the systemU - internal energy of the system,p - pressureV - volume of the system

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Enthalpy (from the Greek enthalpein, “to warm”)is also the state function measured as heattransferred to or from the system:∆H = ∆U + p∆V

- under constant pressure (p) and volume (∆V = 0):∆H = ∆U

∆H - the change of enthalpy (J/mol)∆U - the change of the internal energy

Enthalpy as the state function:

∆H = Hproducts - Hreactants

∆H > 0 reactants are more stable∆H = 0 neither are more stable∆H < 0 products are more stable

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• the universe (the system + surroundings)always tends toward increasing disorder: in allnatural processes, the entropy of the universeincreases.

The second law of thermodynamics

This is why This is why we donwe don’’ t t teach our children about entropyteach our children about entropyuntil until ... (... (youryour ageage?)?)

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• the universe (the system + surroundings)always tends toward increasing disorder: in allnatural processes, the entropy of the universeincreases.

Entropy (S) - is a measure of the disorder ofthe system, it defines the extent of randomnessor disorder of the system.

Any change in randomness of the system isexpressed as the entropy change, ∆S, which byconvention has a positive value when randomnessincreases.

∆S = Sproducts - Sreactants

The second law of thermodynamics

We can say:

The second law of thermodynamics, which can be stated in several forms, is

the law that formulates the spontaneity of processes

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∆Ssystem > 0 it implies that the system becomes more disordered during the reaction and the reaction is spontaneous

∆Ssystem < 0 it implies that the system becomes less disordered (more structurated) during the reaction, the reaction requires energy to occur

The mathematical formula isS = Q/T or ∆S = ∆H/T

∆S – change of entropyQ - the flow of heat to or from the systemT - absolute temperature (K)

∆H – change of enthalpy

The total change of entropy:∆Stotal = ∆Ssystem + ∆Ssurroundings

thermodynamic equilibrium = the state withthe maximal value of entropy

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Living cell is in a low-entropy, nonequilibrium state characterized by a high degree of structural organization.

Cell must release some of the energy it obtains from its environment as heat, thereby increasing Ssurr .

Gibbs free energy& sponaeity of the process

ΔG = ΔH –TΔS

Gibbs free energy expresses the amount of energy capable of doing work during a reaction at constant temperature and pressure.

∆G = Gproducts - Greactants

JosiahWillard Gibbs

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ΔG < 0 → the reaction is spontaneous in theforward direction (exergonic)

ΔG > 0 → the reaction is nonspontaneous inthe forward direction (endergonic)

ΔG = 0 → the system is at equilibrium

ΔG = ΔH –TΔS

Exergonic reaction (energy released, spontaneous)

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Endergonic reaction(energy required, non-spontaneous)

The standard Gibbs free energy ΔG0 for

a reaction refers to the energy change for

a reaction starting at 1 M substrate and product

concentrations and proceeding to equilibrium at

25°C and 101,3 kPa (1 atm).

ΔG´0 - conditions at pH = 7 (standard transformed constant)

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Work and energy consist of two quantities: an intensity factor (driving force), and capacity factor (measure of substance quantity).

ReactantsCurrent

sourceElevated position

The forms of work

Reactants

The forms of work

A system is capable of performing work when matter is moving along a potential gradient.

A molecule is capable of performing chemicalwork when it is moving along a chemicalpotential gradient (change of G)

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In equilibrium ∆G = 0, the equation is transformed to:ΔG'° = - RT ln Keq

R - the universal gas constant (8.31 J/molK), T - absolute temperature (K).

∆G for any chemical reaction = function of thestandard free-energy change, ∆G'0, a constant that ischaracteristic of each specific reaction, and a termthat expresses the initial concentrations ofreactants and products:

Dependence of ∆G on reactant and product concentrations

ΔG'° = - RT ln Keq

Keq = [glu-1-P]/[glu-6-P]

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Keq = [glu-1-P]/[glu-6-P]

ΔG'° = -RT ln [glu-1-P]/[glu-6-P]

+6,890 = -8.31 × 298 ln [glu-1-P]/[glu-6-P] (negative value!)

[glu-1-P]/[glu-6-P] = e-2.78

[glu-1-P]/[glu-6-P] = 0.062

so the ratio of [glu-1-P] to [glu-6-P] at equilibrium is 0.062.

If this ratio P/S will decrease to e.g. 0.032 (because another

reaction uses glu-1-P):

ΔG = ΔG'° + RT ln [glu-1-P]/[glu-6-P]

ΔG = 6,890 + 8.31× 298 ln 0.032

ΔG = - 1,630 J/mol = -1.63 kJ/mol (negative value!)

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Light energy

photosynthesis

Biological workBiological work

Chemical energyChemical energy

RespirationRespiration

ADPADP ATPATPPP

Saccharides and other productsSaccharides and other products

OxygenOxygen

WaterWater

COCO22

6 6 COCO22 + + 6 6 HH22O O CC66HH1212OO66 + O+ O22

chlorophyl h.chlorophyl h.עע

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ATP- plays special role

∆G'0 of ATP hydrolysis = -30.5 kJ/mol (under standard conditions)

Energy charge of the cell = an index used to measure the energy status of biological cells. • related to ATP, ADP and AMP concentrations and it is defined as:

[ATP] + 1/2[ADP][AMP] + [ADP] + [ATP]

↑ AMP or ADP - stimulation of catabolic reactions to produce ATP

.

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The energy of ATP serves for:

Exergonic reaction can be coupled to anendergonic reaction to drive otherwise unfavorablereactions.Preconditions: the ΔG'° values in a reactionsequence are additive, the pathway acquires anoverall negative ΔG'° and there is a commonintermediate in the reaction pathway. Shortly:

The first reaction A → B ΔG1'°

The second reaction B → C ΔG2'°

Sum: A → C ΔG1'° + ΔG2

'°

Coupling of reactions

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Coupling of reactionsSYNTHESIS OF GLUTAMINE

∆ GGlu = +14.3 kJ/mol

Is this reaction catabolic or anabolic?Anabolic

Is this reaction exergonic or endergonic?Endergonic

Will this reaction happen spontaneously?No

Conversion reaction - coupled with ATP hydrolysis

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Free energy change four coupled reaction

+ ∆GATP = -30.5 kJ/mol

∆GGlu = +14.3 kJ/mol

∆GGlu = +14.3 kJ/mol

+ ∆GATP = -30.5 kJ/mol

net ∆G = -16.2 kJ/mol

ENERGY RICH COMPOUNDS

• the high-energy bonds, called energy rich bonds(marked as ~)

•not a special kind of chemical bonds.

•energy rich bonds = covalent bonds - when hydrolyzed, energy is released

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Other energy rich compounds

(Self-study)

∆G´0 = - 61.9 kJ/mol

PhosphoenolpyruvatePhosphoenolpyruvate((enolphopshateenolphopshate energyenergy richrich bondbond))

COOH

C O PCH2

~

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∆G´0 = - 49.3 kJ/mol

1,31,3--bisphosphoglyceratebisphosphoglycerate((acylphopshateacylphopshate energyenergy richrich bondbond))

AcetylCoAAcetylCoA((thioesterthioester energyenergy richrich bondbond))

∆G´0 = - 31.4 kJ/mol

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Acetyl CoA

- source of phosphoryl groups for the quick synthesis ofATP from ADP in the case of extended energy duringmuscle contraction. The enzyme creatine kinase catalyzesthe reversible reaction

ADP + PCr �� ATP + Cr

PhosphocreatinePhosphocreatine((CreatineCreatine phosphatephosphate))

((guanidiniumguanidinium phosphatephosphate energyenergy richrich bondbond))

∆G´0 = - 43.0 kJ/mol

CHEMICAL EQUILIBRIUM EQUILIBRIUM CONSTANT LE ... · 7. 4. 2020 1 •• acids & bases ••...

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