bioenergetics

kinetic energy (4)

mechanical

electrical

sound

thermal

potential energy (3)

gravitational

electrostatic

chemical potential energy

energy

capacity to do work

photosynthesis

6CO2 + 6H2O + energy → C6H12O6 + 6O2

respiration

C6H12O6 + 6O2 → 6CO2 + 6H2O + energy

1st law of thermodynamics (3)

energy can be changed from one form to the other ; it can be neither created nor destroyed

total energy of universe = constant

law of conservation of energy

universe

system + surroundings

open system

can exchange energy + matter with the surroundings

closed system

can exchange energy but NOT matter with the surroundings

isolate system

can exchange neither energy nor matter with the surroundings

endothermic

ΔT increases (+)

q (+)

heat transfer from surr → sys

exothermic

ΔT decreases (-)

q (-)

heat transfer from sys → surr

SI unit for energy

J = joule

quantity of heat transfered depends on

size of temp diff

quantity of material

identity of material losing / gaining heat

specific heat capacity

quantity of heat required to raise the temp. of 1 gram of a substance by one kelvin

Q= mcΔT

Q= heat transferred (J)

m= mass (g)

c= specific heat capacity (J/gK)

ΔT= change in temp (K)

molar heat capacity

at constant pressure (cp)

ΔE = q + w

ΔE= change in internal energy (J)

q= heat transferred to/from system (J)

w= work transferred to/from system (J)

w= -pΔV

w= work done

p=pressure

ΔV= change in volume

ΔE = q-pΔV

ΔE= change in internal energy

q= heat transferred to/from surroundings

p= pressure

ΔV= change in volume

heat

transfer of energy that gives rise to chaotic motion in the surroundings (system)

work

transfer of energy that gives rise to a uniform motion in the surroundings (system)

heat (surr→ sys)

q>0 (+)

E increases

heat (sys→surr)

q<0 (-)

E decreases

work (surr→sys)

w>0 (+)

E increases

work (sys→ surr)

w<0 (-)

E decreases

internal energy of chemical sys=

the sum of kinetic and potential energies of atoms, mols / ions in sys

kinetic energy

energy of motion of atoms, molecules or ions

larger heat capacity, c=

more heat stored

c of solids liquids and gases

solids and liquids> gases

more complex molecules c

more heat stored

potential energy

attractive/repulsive force between nulcei + electrons in the system

intra vs intermolecular

intra> inter

breaking/ making bonds

break= requires energy

make= releases energy

enthalpy, H

heat content of a substance at constant pressure

enthalpy change ΔH

heat absorbed/ released by the system at constant pressure

hess’s law

if a reaction is the sum (difference) of 2 or more reactions, ΔrH of the overall reaction is the sum (difference) of the ΔrH values of these reactions

standard state of an element or compound

defined as the most stable form of the substance in the state that exists a pressure of 1 bar + specified temp.

standard enthalpy of reaction ΔrH⦵

enthalpy of a reaction which converts pure reactants in their standard states into pure products in their standard states

standard enthalpy of formation ΔfH⦵

enthalpy of a reaction in which a substance is made directly from its elements in their standard states

most are -ve (exo)

standard enthalpy of formation of products - reactants = standard enthalpy of reaction

ΔrH⦵ = ∑ΔfH⦵(products) - ∑ΔfH⦵(reactants)

standard enthalpy of combustion ΔcH⦵

enthalpy change when one mole of a substance is totally combusted in oxygen at 298K and 1atm

bomb calorimeter

sample ignited in chamber containing oxygen at high pressure

combustion chamber in water in well insulated outer chamber (heat from combustion passes to water)

can relate enthalpy change to increase in temp of water

standard enthalpy of combustion used to calc standard enthalpy of formation using hess’ law

ΔfH⦵ = ∑ΔcH⦵ (rea) - ∑ΔcH⦵ (pro)

exothermic reactions

(-ΔrH⦵) transfer energy to the surroundings, generally product favoured (spontaneous)

endothermic reactions

(+ΔrH⦵) transfer energy from the surroundings, generally reactant favoured (non-spon)

spontaneous

natural tendency to occur (ready for change)

non-spontaneous

requires work to be done on the system for the change to be brought about

lower disorder=

increased order

final state is more probable than its initial state if:

both energy + matter are dispersed

if only energy or matter is dispersed a quantitative consideration is needed to decide if the process is spontaneous or not

2nd law of thermodynamics

in a spontaneous process the entropy of the universe increases

entropy

thermodynamic measure of the tendency for the universe to move towards an increased disorder (S)

third law of thermodynamics

at absolute zero, entropy of a perfectly crystalline solid is zero

above 0K, all substances have S>0

standard molar entropy

the entropy gained by converting one mole from a perfect crystal at 0K to standard state conditions

in general g>l>s

larger and more complicated molecules have bigger sme

standard enthalpy of reactions, ΔrS⦵

ΔrS⦵ = ∑S⦵ (pro) - ∑S⦵ (rea)

gibbs free energy

criterion of spontaneity

ΔrG<0

spontaneous reaction

ΔrG>0

non-spon reaction

ΔrG=0

reaction at equilibrium

standard free energy of reaction calc

ΔrG⦵= ΔrH⦵ - ΔrS⦵T

standard free energy of formation

ΔfG⦵ gibbs free energy of formation of one mole of a compound from component elements in their standard states

calculating ΔrG⦵ from formation values

ΔrG⦵ = ΣΔfG⦵ (pro) - ΣΔfG⦵ (rea)

ΔG

max work available from a process

pro vs rea gibbs energy

products have lower gibbs

H= -ve

S= +ve

G= -ve

spontaneous

H= -ve

S= -ve

G= depends

spon at lower temp

H= +ve

S= +ve

G= depends

spon at higher temp

H= +ve

S= -ve

G= +ve

non spon

metabolic pathway

series of reactions that overall are energetically favourable as a result of sequential and/or parallel coupling

2 types of metabolic reactions

catabolic and anabolic

catabolic

break down complex molecules to simpler ones (exergonic- release gibbs)

anabolic

synthesise complex molecules from simpler ones (endergonic= requires energy)

coupled ana and cata reactions

energy released by catabolic are used to drive anabolic ones

ATP-ADP cycle

ATP

dephosphorylation (-/spon)

ADP

phosphorylation (+/non-spon)

oxidation of glucose (resp)

exergonic (-)

glycolysis= 1 molecule of glucose → 2 molecules of pyruvate

2 ATP molecules formed in process

how many moles of ATP synthesised per mole of glucose oxidised

32-36 moles

parallel coupling of chemical reactions

simultaneous occurence of unfavourable + favourable reactions

phospholipid bilayer

polar head

non-polar tail

formation of lipid bilayers spon or non-spon

spon

folded protein

more constrained (intramolecular forces) reduces disorder

hydrophobic groups in protein

go to inside/ middle of protein

diff. between ΔG and ΔH

energy which will be used to create a more organised system and must be discarded as entropy