Bioenergetics looks at how energy moves through and changes form within living systems. This idea sits behind many cellular processes, including metabolism and biosynthesis.

Thermodynamics gives us a framework for predicting whether a reaction is energetically favorable by tracking enthalpy (H), entropy (S), and free energy (G). The key relationship is:

$Δ G = Δ H - T Δ S$

A process is spontaneous when $Δ G$ is negative and nonspontaneous when $Δ G$ is positive.

Free energy/ $K_{e q}$

The equilibrium constant ( $K_{e q}$ ) describes the ratio of products to reactants at equilibrium. Under standard conditions, free energy and equilibrium are related by:

$Δ G^{\circ} = - RT ln (K_{e q})$

A large $K_{e q}$ means products are favored at equilibrium, but it does not mean the reaction happens quickly. Reaction rate is a kinetics question, not a thermodynamics one.

Concentration

Changing the concentration of reactants or products shifts equilibrium according to Le Châtelier’s Principle. By raising or lowering the amount of a species in a reaction, cells can push the reaction forward or backward to help maintain metabolic balance.

Phosphorylation/ATP

In cells, ATP (adenosine triphosphate) serves as a readily available energy currency. Phosphorylation transfers a phosphate group to a substrate or enzyme, often changing its reactivity or function.

ATP hydrolysis $Δ G < 0$
- Breaking ATP into ADP + Pi has a strongly negative $Δ G$ . This exergonic energy release can be coupled to endergonic processes such as muscle contraction and macromolecule synthesis.
ATP group transfers
- ATP can also transfer its phosphate group to form phosphorylated intermediates. These “activated” compounds tend to react more readily, which helps drive biosynthetic pathways.

Biological oxidation-reduction

Organisms often capture energy through redox reactions, where electrons move from nutrients to electron acceptors. The energy released can be used to support ATP production and other endergonic steps in metabolism.

Half-reactions Redox reactions can be split into two half-reactions: oxidation (loss of electrons) and reduction (gain of electrons). Writing half-reactions makes it easier to track electron flow and connect that flow to energy changes.
Soluble electron carriers Molecules such as $N A D^{+}$ , $F A D$ , and ubiquinone carry electrons between substrates and the electron transport chain. They cycle between oxidized and reduced forms, allowing energy to be released in steps rather than all at once.
Flavoproteins Some redox enzymes contain flavin groups ( $FMN$ or $F A D$ ) as prosthetic groups. These flavoproteins are key players in electron transport because the flavin ring can repeatedly accept and donate electrons.

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Chemical reactions can either absorb heat from the surroundings or release heat to them.

Endothermic reactions absorb heat ( $Δ H > 0$ ).
Exothermic reactions release heat ( $Δ H < 0$ ).

Standard heats of reaction and formation quantify these heat changes under specified (standard) conditions and tell you whether heat is required or liberated during the transformation.

Enthalpy and standard heats of reaction

Enthalpy (H) is the heat content of a system at constant pressure. The enthalpy change ( $Δ H$ ) for a reaction is often reported under standard conditions and is typically given in joules per mole.

Standard heat of reaction ( $Δ H_{r x n}$ ): the net enthalpy change when reactants are converted to products.
Standard heat of formation ( $Δ H_{f}$ ): the enthalpy change when 1 mole of a compound forms from its elements in their standard states.

Hess’s law

Hess’s Law says the total enthalpy change for a reaction does not depend on the path taken to get from reactants to products. That means you can find $Δ H$ for an overall reaction by adding the $Δ H$ values of steps that sum to the same net reaction:

$Δ H_{rxn} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

Whether a reaction is thermodynamically favorable depends on its free energy change ( $Δ G$ ).

If $Δ G < 0$ , the reaction is thermodynamically favorable (spontaneous) under the given conditions.
If $Δ G > 0$ , the reaction is thermodynamically unfavorable (nonspontaneous) under the given conditions.

Exothermic reactions (negative $Δ H$ ) often have negative $Δ G$ , but $Δ H$ alone doesn’t decide spontaneity. Entropy matters too, through:

$Δ G = Δ H - T Δ S$

Phosphoryl group transfers and ATP

In biological systems, energy transfer is often driven by ATP hydrolysis. ATP hydrolysis is highly exergonic ( $Δ G << 0$ ), and cells can couple that energy release to phosphoryl group transfer reactions. These transfers help activate molecules, regulate metabolic pathways, and power many cellular functions.

Biological oxidation-reduction

Cellular energy production relies on oxidation-reduction (redox) reactions, where electrons are transferred between molecules. These reactions can be represented as half-reactions and are often mediated by soluble electron carriers such as $N A D^{+}$ and $F A D$ . Flavoproteins and other enzymes facilitate these electron transfers and are central to bioenergetic pathways that convert chemical energy into usable forms.

Thermodynamic system - state function

A thermodynamic system is described using state functions - properties such as temperature, pressure, volume, and internal energy. A state function depends only on the system’s current state, not on the path taken to reach that state.

Zeroth law - concept of temperature

The zeroth law states that if two systems are each in thermal equilibrium with a third system, then they are in thermal equilibrium with each other. This principle underlies the concept of temperature and explains why temperature determines the direction of heat exchange.

First law - conservation of energy in thermodynamic processes

The first law expresses conservation of energy: the change in a system’s internal energy equals the heat added to the system minus the work done by the system. Energy isn’t created or destroyed; it is converted from one form to another.

PV diagram: work done = area under or enclosed by curve (PHY)

On a pressure-volume (PV) diagram, the work done by a system during expansion or compression corresponds to the area under (or enclosed by) the process curve. This makes work easy to visualize as a geometric area on the graph.

Isothermal process in ideal gas on a p-V diagram

Second law - concept of entropy

The second law introduces entropy as a measure of randomness or dispersal of energy. It states that in any natural process, the total entropy of an isolated system tends to increase. This tendency explains why many spontaneous changes are irreversible.

Entropy as a measure of “disorder”

Entropy describes molecular disorder. In general, gases have high entropy because molecules move freely, liquids have moderate entropy, and crystalline solids have low entropy because their particles are arranged in an ordered structure.

Relative entropy for gas, liquid, and crystal states

States of matter differ in entropy:

Gases have the highest entropy because molecules are widely dispersed.
Liquids have intermediate entropy.
Crystals are highly ordered and have the lowest entropy.

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry measures heat changes during reactions or physical processes.

Heat capacity: the heat required to raise the temperature of a system.
Specific heat: the heat required per unit mass to raise temperature by one degree.

Heat transfer - conduction, convection, radiation

Heat moves in three main ways:

Conduction: transfer through direct molecular collisions.
Convection: transfer through bulk motion of a fluid.
Radiation: transfer through electromagnetic waves.

Bond dissociation energy as related to heats of formation

Bond dissociation energy is the energy required to break a chemical bond. It relates to heats of formation because bond strengths influence how stable molecules are and therefore affect the overall energy change during chemical reactions.

Spontaneous reactions and $Δ G °$

Spontaneity under standard conditions is described by the standard free energy change, $Δ G °$ .

If $Δ G ° < 0$ , the reaction is spontaneous under standard conditions.
If $Δ G ° > 0$ , the reaction is nonspontaneous under standard conditions.

Coefficient of expansion

The coefficient of expansion describes how a material’s volume or length changes with temperature. Materials with larger coefficients expand more when heated, which matters in both engineering contexts and thermodynamic calculations.

Heat of fusion, heat of vaporization

Heat of fusion: the energy required to convert a solid to a liquid at its melting point.
Heat of vaporization: the energy required to convert a liquid to a gas at its boiling point.

Both quantities reflect the energy needed to overcome intermolecular forces during phase changes.

Phase diagram: pressure and temperature

A phase diagram shows which phase of a substance is stable at different pressures and temperatures. It also marks key points such as the triple point (solid, liquid, and gas coexist) and the critical point (the liquid-gas distinction disappears).

Bioenergetics/thermodynamics

Energy flow and transformations in living systems
Thermodynamics predicts reaction favorability using enthalpy (H), entropy (S), and free energy (G)
$Δ G = Δ H - T Δ S$ ; spontaneous if $Δ G < 0$

Free energy/ $K_{e q}$

$K_{e q}$ : ratio of products to reactants at equilibrium
$Δ G^{\circ} = - RT ln (K_{e q})$
Large $K_{e q}$ favors products; does not indicate reaction rate

Concentration

Le Châtelier’s Principle: changing concentrations shifts equilibrium
Cells adjust concentrations to regulate metabolic pathways

Phosphorylation/ATP

ATP: main cellular energy currency
ATP hydrolysis ( $Δ G < 0$ ) drives endergonic processes
Phosphorylation activates substrates, aiding biosynthesis

Biological oxidation-reduction

Energy captured via redox reactions (electron transfer)
Half-reactions: oxidation (loss), reduction (gain) of electrons
Electron carriers: NAD⁺, FAD, ubiquinone; flavoproteins use FMN/FAD

Endothermic and exothermic reactions

Endothermic: absorb heat ( $Δ H > 0$ )
Exothermic: release heat ( $Δ H < 0$ )
Standard heats quantify energy change under standard conditions

Enthalpy and standard heats of reaction

Enthalpy (H): heat content at constant pressure
$Δ H_{r x n}$ : net enthalpy change for a reaction
$Δ H_{f}$ : enthalpy change for forming 1 mole from elements

Hess’s law

Total enthalpy change independent of reaction path
$Δ H_{r x n} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

$Δ G < 0$ : spontaneous; $Δ G > 0$ : nonspontaneous
Both enthalpy and entropy determine spontaneity ( $Δ G = Δ H - T Δ S$ )

Phosphoryl group transfers and ATP

ATP hydrolysis is highly exergonic ( $Δ G << 0$ )
Drives phosphoryl group transfers, activating molecules and pathways

Biological oxidation-reduction

Redox reactions central to energy production
Mediated by electron carriers (NAD⁺, FAD), flavoproteins

Thermodynamic system - state function

State functions: properties depend only on current state (not path)
Examples: temperature, pressure, volume, internal energy

Zeroth law - concept of temperature

If two systems are in equilibrium with a third, they are in equilibrium with each other
Basis for defining temperature

First law - conservation of energy in thermodynamic processes

Energy conserved: $Δ U = Q - W$
- $Δ U$ : change in internal energy
- $Q$ : heat added, $W$ : work done by system

PV diagram: work done = area under or enclosed by curve

Work visualized as area under process curve on PV diagram

Second law - concept of entropy

Entropy: measure of randomness/disorder
Total entropy of isolated system increases in natural processes

Entropy as a measure of “disorder”

Gases: high entropy; liquids: intermediate; crystals: low entropy

Relative entropy for gas, liquid, and crystal states

Gases > liquids > crystals (in entropy)

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry: measures heat changes
Heat capacity: heat needed to raise system temperature
Specific heat: heat per unit mass per degree

Heat transfer - conduction, convection, radiation

Conduction: direct contact
Convection: fluid movement
Radiation: electromagnetic waves

Bond dissociation energy as related to heats of formation

Bond dissociation energy: energy to break a bond
Influences heats of formation and reaction energetics

Spontaneous reactions and $Δ G^{\circ}$

$Δ G^{\circ} < 0$ : spontaneous under standard conditions
$Δ G^{\circ} > 0$ : nonspontaneous under standard conditions

Coefficient of expansion

Describes how material size changes with temperature
Larger coefficient = more expansion when heated

Heat of fusion, heat of vaporization

Heat of fusion: energy to melt solid at melting point
Heat of vaporization: energy to boil liquid at boiling point

Phase diagram: pressure and temperature

Shows stable phases at various pressures and temperatures
Key points: triple point, critical point

Bioenergetics, thermochemistry and thermodynamics

Bioenergetics/thermodynamics

Bioenergetics looks at how energy moves through and changes form within living systems. This idea sits behind many cellular processes, including metabolism and biosynthesis.

Thermodynamics gives us a framework for predicting whether a reaction is energetically favorable by tracking enthalpy (H), entropy (S), and free energy (G). The key relationship is:

$Δ G = Δ H - T Δ S$

A process is spontaneous when $Δ G$ is negative and nonspontaneous when $Δ G$ is positive.

Free energy/ $K_{e q}$

The equilibrium constant ( $K_{e q}$ ) describes the ratio of products to reactants at equilibrium. Under standard conditions, free energy and equilibrium are related by:

$Δ G^{\circ} = - RT ln (K_{e q})$

A large $K_{e q}$ means products are favored at equilibrium, but it does not mean the reaction happens quickly. Reaction rate is a kinetics question, not a thermodynamics one.

Concentration

Phosphorylation/ATP

ATP hydrolysis $Δ G < 0$
- Breaking ATP into ADP + Pi has a strongly negative $Δ G$ . This exergonic energy release can be coupled to endergonic processes such as muscle contraction and macromolecule synthesis.
ATP group transfers
- ATP can also transfer its phosphate group to form phosphorylated intermediates. These “activated” compounds tend to react more readily, which helps drive biosynthetic pathways.

Biological oxidation-reduction

Half-reactions Redox reactions can be split into two half-reactions: oxidation (loss of electrons) and reduction (gain of electrons). Writing half-reactions makes it easier to track electron flow and connect that flow to energy changes.
Soluble electron carriers Molecules such as $N A D^{+}$ , $F A D$ , and ubiquinone carry electrons between substrates and the electron transport chain. They cycle between oxidized and reduced forms, allowing energy to be released in steps rather than all at once.
Flavoproteins Some redox enzymes contain flavin groups ( $FMN$ or $F A D$ ) as prosthetic groups. These flavoproteins are key players in electron transport because the flavin ring can repeatedly accept and donate electrons.

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Chemical reactions can either absorb heat from the surroundings or release heat to them.

Endothermic reactions absorb heat ( $Δ H > 0$ ).
Exothermic reactions release heat ( $Δ H < 0$ ).

Standard heats of reaction and formation quantify these heat changes under specified (standard) conditions and tell you whether heat is required or liberated during the transformation.

Enthalpy and standard heats of reaction

Standard heat of reaction ( $Δ H_{r x n}$ ): the net enthalpy change when reactants are converted to products.
Standard heat of formation ( $Δ H_{f}$ ): the enthalpy change when 1 mole of a compound forms from its elements in their standard states.

Hess’s law

$Δ H_{rxn} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

Whether a reaction is thermodynamically favorable depends on its free energy change ( $Δ G$ ).

If $Δ G < 0$ , the reaction is thermodynamically favorable (spontaneous) under the given conditions.
If $Δ G > 0$ , the reaction is thermodynamically unfavorable (nonspontaneous) under the given conditions.

Exothermic reactions (negative $Δ H$ ) often have negative $Δ G$ , but $Δ H$ alone doesn’t decide spontaneity. Entropy matters too, through:

$Δ G = Δ H - T Δ S$

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

States of matter differ in entropy:

Gases have the highest entropy because molecules are widely dispersed.
Liquids have intermediate entropy.
Crystals are highly ordered and have the lowest entropy.

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry measures heat changes during reactions or physical processes.

Heat capacity: the heat required to raise the temperature of a system.
Specific heat: the heat required per unit mass to raise temperature by one degree.

Heat transfer - conduction, convection, radiation

Heat moves in three main ways:

Conduction: transfer through direct molecular collisions.
Convection: transfer through bulk motion of a fluid.
Radiation: transfer through electromagnetic waves.

Bond dissociation energy as related to heats of formation

Spontaneous reactions and $Δ G °$

Spontaneity under standard conditions is described by the standard free energy change, $Δ G °$ .

If $Δ G ° < 0$ , the reaction is spontaneous under standard conditions.
If $Δ G ° > 0$ , the reaction is nonspontaneous under standard conditions.

Coefficient of expansion

Heat of fusion, heat of vaporization

Heat of fusion: the energy required to convert a solid to a liquid at its melting point.
Heat of vaporization: the energy required to convert a liquid to a gas at its boiling point.

Both quantities reflect the energy needed to overcome intermolecular forces during phase changes.

Phase diagram: pressure and temperature

Achievable MCAT

4. Chem/phys

4.10. Principles of chemical thermodynamics and kinetics, enzymes

Bioenergetics, thermochemistry and thermodynamics

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Bioenergetics/thermodynamics

Bioenergetics looks at how energy moves through and changes form within living systems. This idea sits behind many cellular processes, including metabolism and biosynthesis.

Thermodynamics gives us a framework for predicting whether a reaction is energetically favorable by tracking enthalpy (H), entropy (S), and free energy (G). The key relationship is:

$Δ G = Δ H - T Δ S$

A process is spontaneous when $Δ G$ is negative and nonspontaneous when $Δ G$ is positive.

Free energy/ $K_{e q}$

The equilibrium constant ( $K_{e q}$ ) describes the ratio of products to reactants at equilibrium. Under standard conditions, free energy and equilibrium are related by:

$Δ G^{\circ} = - RT ln (K_{e q})$

A large $K_{e q}$ means products are favored at equilibrium, but it does not mean the reaction happens quickly. Reaction rate is a kinetics question, not a thermodynamics one.

Concentration

Phosphorylation/ATP

ATP hydrolysis $Δ G < 0$
- Breaking ATP into ADP + Pi has a strongly negative $Δ G$ . This exergonic energy release can be coupled to endergonic processes such as muscle contraction and macromolecule synthesis.
ATP group transfers
- ATP can also transfer its phosphate group to form phosphorylated intermediates. These “activated” compounds tend to react more readily, which helps drive biosynthetic pathways.

Biological oxidation-reduction

Half-reactions Redox reactions can be split into two half-reactions: oxidation (loss of electrons) and reduction (gain of electrons). Writing half-reactions makes it easier to track electron flow and connect that flow to energy changes.
Soluble electron carriers Molecules such as $N A D^{+}$ , $F A D$ , and ubiquinone carry electrons between substrates and the electron transport chain. They cycle between oxidized and reduced forms, allowing energy to be released in steps rather than all at once.
Flavoproteins Some redox enzymes contain flavin groups ( $FMN$ or $F A D$ ) as prosthetic groups. These flavoproteins are key players in electron transport because the flavin ring can repeatedly accept and donate electrons.

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Chemical reactions can either absorb heat from the surroundings or release heat to them.

Endothermic reactions absorb heat ( $Δ H > 0$ ).
Exothermic reactions release heat ( $Δ H < 0$ ).

Standard heats of reaction and formation quantify these heat changes under specified (standard) conditions and tell you whether heat is required or liberated during the transformation.

Enthalpy and standard heats of reaction

Standard heat of reaction ( $Δ H_{r x n}$ ): the net enthalpy change when reactants are converted to products.
Standard heat of formation ( $Δ H_{f}$ ): the enthalpy change when 1 mole of a compound forms from its elements in their standard states.

Hess’s law

$Δ H_{rxn} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

Whether a reaction is thermodynamically favorable depends on its free energy change ( $Δ G$ ).

If $Δ G < 0$ , the reaction is thermodynamically favorable (spontaneous) under the given conditions.
If $Δ G > 0$ , the reaction is thermodynamically unfavorable (nonspontaneous) under the given conditions.

Exothermic reactions (negative $Δ H$ ) often have negative $Δ G$ , but $Δ H$ alone doesn’t decide spontaneity. Entropy matters too, through:

$Δ G = Δ H - T Δ S$

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

States of matter differ in entropy:

Gases have the highest entropy because molecules are widely dispersed.
Liquids have intermediate entropy.
Crystals are highly ordered and have the lowest entropy.

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry measures heat changes during reactions or physical processes.

Heat capacity: the heat required to raise the temperature of a system.
Specific heat: the heat required per unit mass to raise temperature by one degree.

Heat transfer - conduction, convection, radiation

Heat moves in three main ways:

Conduction: transfer through direct molecular collisions.
Convection: transfer through bulk motion of a fluid.
Radiation: transfer through electromagnetic waves.

Bond dissociation energy as related to heats of formation

Spontaneous reactions and $Δ G °$

Spontaneity under standard conditions is described by the standard free energy change, $Δ G °$ .

If $Δ G ° < 0$ , the reaction is spontaneous under standard conditions.
If $Δ G ° > 0$ , the reaction is nonspontaneous under standard conditions.

Coefficient of expansion

Heat of fusion, heat of vaporization

Heat of fusion: the energy required to convert a solid to a liquid at its melting point.
Heat of vaporization: the energy required to convert a liquid to a gas at its boiling point.

Both quantities reflect the energy needed to overcome intermolecular forces during phase changes.

Phase diagram: pressure and temperature

Bioenergetics/thermodynamics

Energy flow and transformations in living systems
Thermodynamics predicts reaction favorability using enthalpy (H), entropy (S), and free energy (G)
$Δ G = Δ H - T Δ S$ ; spontaneous if $Δ G < 0$

Free energy/ $K_{e q}$

$K_{e q}$ : ratio of products to reactants at equilibrium
$Δ G^{\circ} = - RT ln (K_{e q})$
Large $K_{e q}$ favors products; does not indicate reaction rate

Concentration

Le Châtelier’s Principle: changing concentrations shifts equilibrium
Cells adjust concentrations to regulate metabolic pathways

Phosphorylation/ATP

ATP: main cellular energy currency
ATP hydrolysis ( $Δ G < 0$ ) drives endergonic processes
Phosphorylation activates substrates, aiding biosynthesis

Biological oxidation-reduction

Energy captured via redox reactions (electron transfer)
Half-reactions: oxidation (loss), reduction (gain) of electrons
Electron carriers: NAD⁺, FAD, ubiquinone; flavoproteins use FMN/FAD

Endothermic and exothermic reactions

Endothermic: absorb heat ( $Δ H > 0$ )
Exothermic: release heat ( $Δ H < 0$ )
Standard heats quantify energy change under standard conditions

Enthalpy and standard heats of reaction

Enthalpy (H): heat content at constant pressure
$Δ H_{r x n}$ : net enthalpy change for a reaction
$Δ H_{f}$ : enthalpy change for forming 1 mole from elements

Hess’s law

Total enthalpy change independent of reaction path
$Δ H_{r x n} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

$Δ G < 0$ : spontaneous; $Δ G > 0$ : nonspontaneous
Both enthalpy and entropy determine spontaneity ( $Δ G = Δ H - T Δ S$ )

Phosphoryl group transfers and ATP

ATP hydrolysis is highly exergonic ( $Δ G << 0$ )
Drives phosphoryl group transfers, activating molecules and pathways

Biological oxidation-reduction

Redox reactions central to energy production
Mediated by electron carriers (NAD⁺, FAD), flavoproteins

Thermodynamic system - state function

State functions: properties depend only on current state (not path)
Examples: temperature, pressure, volume, internal energy

Zeroth law - concept of temperature

If two systems are in equilibrium with a third, they are in equilibrium with each other
Basis for defining temperature

First law - conservation of energy in thermodynamic processes

Energy conserved: $Δ U = Q - W$
- $Δ U$ : change in internal energy
- $Q$ : heat added, $W$ : work done by system

PV diagram: work done = area under or enclosed by curve

Work visualized as area under process curve on PV diagram

Second law - concept of entropy

Entropy: measure of randomness/disorder
Total entropy of isolated system increases in natural processes

Entropy as a measure of “disorder”

Gases: high entropy; liquids: intermediate; crystals: low entropy

Relative entropy for gas, liquid, and crystal states

Gases > liquids > crystals (in entropy)

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry: measures heat changes
Heat capacity: heat needed to raise system temperature
Specific heat: heat per unit mass per degree

Heat transfer - conduction, convection, radiation

Conduction: direct contact
Convection: fluid movement
Radiation: electromagnetic waves

Bond dissociation energy as related to heats of formation

Bond dissociation energy: energy to break a bond
Influences heats of formation and reaction energetics

Spontaneous reactions and $Δ G^{\circ}$

$Δ G^{\circ} < 0$ : spontaneous under standard conditions
$Δ G^{\circ} > 0$ : nonspontaneous under standard conditions

Coefficient of expansion

Describes how material size changes with temperature
Larger coefficient = more expansion when heated

Heat of fusion, heat of vaporization

Heat of fusion: energy to melt solid at melting point
Heat of vaporization: energy to boil liquid at boiling point

Phase diagram: pressure and temperature

Shows stable phases at various pressures and temperatures
Key points: triple point, critical point

Bioenergetics, thermochemistry and thermodynamics

Bioenergetics/thermodynamics

Bioenergetics looks at how energy moves through and changes form within living systems. This idea sits behind many cellular processes, including metabolism and biosynthesis.

Thermodynamics gives us a framework for predicting whether a reaction is energetically favorable by tracking enthalpy (H), entropy (S), and free energy (G). The key relationship is:

$Δ G = Δ H - T Δ S$

A process is spontaneous when $Δ G$ is negative and nonspontaneous when $Δ G$ is positive.

Free energy/ $K_{e q}$

The equilibrium constant ( $K_{e q}$ ) describes the ratio of products to reactants at equilibrium. Under standard conditions, free energy and equilibrium are related by:

$Δ G^{\circ} = - RT ln (K_{e q})$

A large $K_{e q}$ means products are favored at equilibrium, but it does not mean the reaction happens quickly. Reaction rate is a kinetics question, not a thermodynamics one.

Concentration

Phosphorylation/ATP

ATP hydrolysis $Δ G < 0$
- Breaking ATP into ADP + Pi has a strongly negative $Δ G$ . This exergonic energy release can be coupled to endergonic processes such as muscle contraction and macromolecule synthesis.
ATP group transfers
- ATP can also transfer its phosphate group to form phosphorylated intermediates. These “activated” compounds tend to react more readily, which helps drive biosynthetic pathways.

Biological oxidation-reduction

Half-reactions Redox reactions can be split into two half-reactions: oxidation (loss of electrons) and reduction (gain of electrons). Writing half-reactions makes it easier to track electron flow and connect that flow to energy changes.
Soluble electron carriers Molecules such as $N A D^{+}$ , $F A D$ , and ubiquinone carry electrons between substrates and the electron transport chain. They cycle between oxidized and reduced forms, allowing energy to be released in steps rather than all at once.
Flavoproteins Some redox enzymes contain flavin groups ( $FMN$ or $F A D$ ) as prosthetic groups. These flavoproteins are key players in electron transport because the flavin ring can repeatedly accept and donate electrons.

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Chemical reactions can either absorb heat from the surroundings or release heat to them.

Endothermic reactions absorb heat ( $Δ H > 0$ ).
Exothermic reactions release heat ( $Δ H < 0$ ).

Standard heats of reaction and formation quantify these heat changes under specified (standard) conditions and tell you whether heat is required or liberated during the transformation.

Enthalpy and standard heats of reaction

Standard heat of reaction ( $Δ H_{r x n}$ ): the net enthalpy change when reactants are converted to products.
Standard heat of formation ( $Δ H_{f}$ ): the enthalpy change when 1 mole of a compound forms from its elements in their standard states.

Hess’s law

$Δ H_{rxn} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

Whether a reaction is thermodynamically favorable depends on its free energy change ( $Δ G$ ).

If $Δ G < 0$ , the reaction is thermodynamically favorable (spontaneous) under the given conditions.
If $Δ G > 0$ , the reaction is thermodynamically unfavorable (nonspontaneous) under the given conditions.

Exothermic reactions (negative $Δ H$ ) often have negative $Δ G$ , but $Δ H$ alone doesn’t decide spontaneity. Entropy matters too, through:

$Δ G = Δ H - T Δ S$

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

States of matter differ in entropy:

Gases have the highest entropy because molecules are widely dispersed.
Liquids have intermediate entropy.
Crystals are highly ordered and have the lowest entropy.

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry measures heat changes during reactions or physical processes.

Heat capacity: the heat required to raise the temperature of a system.
Specific heat: the heat required per unit mass to raise temperature by one degree.

Heat transfer - conduction, convection, radiation

Heat moves in three main ways:

Conduction: transfer through direct molecular collisions.
Convection: transfer through bulk motion of a fluid.
Radiation: transfer through electromagnetic waves.

Bond dissociation energy as related to heats of formation

Spontaneous reactions and $Δ G °$

Spontaneity under standard conditions is described by the standard free energy change, $Δ G °$ .

If $Δ G ° < 0$ , the reaction is spontaneous under standard conditions.
If $Δ G ° > 0$ , the reaction is nonspontaneous under standard conditions.

Coefficient of expansion

Heat of fusion, heat of vaporization

Heat of fusion: the energy required to convert a solid to a liquid at its melting point.
Heat of vaporization: the energy required to convert a liquid to a gas at its boiling point.

Both quantities reflect the energy needed to overcome intermolecular forces during phase changes.

Phase diagram: pressure and temperature

Key points

Bioenergetics/thermodynamics

Energy flow and transformations in living systems
Thermodynamics predicts reaction favorability using enthalpy (H), entropy (S), and free energy (G)
$Δ G = Δ H - T Δ S$ ; spontaneous if $Δ G < 0$

Free energy/ $K_{e q}$

$K_{e q}$ : ratio of products to reactants at equilibrium
$Δ G^{\circ} = - RT ln (K_{e q})$
Large $K_{e q}$ favors products; does not indicate reaction rate

Concentration

Le Châtelier’s Principle: changing concentrations shifts equilibrium
Cells adjust concentrations to regulate metabolic pathways

Phosphorylation/ATP

ATP: main cellular energy currency
ATP hydrolysis ( $Δ G < 0$ ) drives endergonic processes
Phosphorylation activates substrates, aiding biosynthesis

Biological oxidation-reduction

Energy captured via redox reactions (electron transfer)
Half-reactions: oxidation (loss), reduction (gain) of electrons
Electron carriers: NAD⁺, FAD, ubiquinone; flavoproteins use FMN/FAD

Endothermic and exothermic reactions

Endothermic: absorb heat ( $Δ H > 0$ )
Exothermic: release heat ( $Δ H < 0$ )
Standard heats quantify energy change under standard conditions

Enthalpy and standard heats of reaction

Enthalpy (H): heat content at constant pressure
$Δ H_{r x n}$ : net enthalpy change for a reaction
$Δ H_{f}$ : enthalpy change for forming 1 mole from elements

Hess’s law

Total enthalpy change independent of reaction path
$Δ H_{r x n} = \sum Δ H_{f} (products) - \sum Δ H_{f} (reactants)$

Spontaneous reactions and standard free energy change

$Δ G < 0$ : spontaneous; $Δ G > 0$ : nonspontaneous
Both enthalpy and entropy determine spontaneity ( $Δ G = Δ H - T Δ S$ )

Phosphoryl group transfers and ATP

ATP hydrolysis is highly exergonic ( $Δ G << 0$ )
Drives phosphoryl group transfers, activating molecules and pathways

Biological oxidation-reduction

Redox reactions central to energy production
Mediated by electron carriers (NAD⁺, FAD), flavoproteins

Thermodynamic system - state function

State functions: properties depend only on current state (not path)
Examples: temperature, pressure, volume, internal energy

Zeroth law - concept of temperature

If two systems are in equilibrium with a third, they are in equilibrium with each other
Basis for defining temperature

First law - conservation of energy in thermodynamic processes

Energy conserved: $Δ U = Q - W$
- $Δ U$ : change in internal energy
- $Q$ : heat added, $W$ : work done by system

PV diagram: work done = area under or enclosed by curve

Work visualized as area under process curve on PV diagram

Second law - concept of entropy

Entropy: measure of randomness/disorder
Total entropy of isolated system increases in natural processes

Entropy as a measure of “disorder”

Gases: high entropy; liquids: intermediate; crystals: low entropy

Relative entropy for gas, liquid, and crystal states

Gases > liquids > crystals (in entropy)

Measurement of heat changes (calorimetry), heat capacity, specific heat

Calorimetry: measures heat changes
Heat capacity: heat needed to raise system temperature
Specific heat: heat per unit mass per degree

Heat transfer - conduction, convection, radiation

Conduction: direct contact
Convection: fluid movement
Radiation: electromagnetic waves

Bond dissociation energy as related to heats of formation

Bond dissociation energy: energy to break a bond
Influences heats of formation and reaction energetics

Spontaneous reactions and $Δ G^{\circ}$

$Δ G^{\circ} < 0$ : spontaneous under standard conditions
$Δ G^{\circ} > 0$ : nonspontaneous under standard conditions

Coefficient of expansion

Describes how material size changes with temperature
Larger coefficient = more expansion when heated

Heat of fusion, heat of vaporization

Heat of fusion: energy to melt solid at melting point
Heat of vaporization: energy to boil liquid at boiling point

Phase diagram: pressure and temperature

Shows stable phases at various pressures and temperatures
Key points: triple point, critical point

Bioenergetics, thermochemistry and thermodynamics

Bioenergetics/thermodynamics

Free energy/Keq​

Concentration

Phosphorylation/ATP

Biological oxidation-reduction

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Enthalpy and standard heats of reaction

Hess’s law

Spontaneous reactions and standard free energy change

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

Measurement of heat changes (calorimetry), heat capacity, specific heat

Heat transfer - conduction, convection, radiation

Bond dissociation energy as related to heats of formation

Spontaneous reactions and ΔG°

Coefficient of expansion

Heat of fusion, heat of vaporization

Phase diagram: pressure and temperature

Bioenergetics, thermochemistry and thermodynamics

Bioenergetics/thermodynamics

Free energy/Keq​

Concentration

Phosphorylation/ATP

Biological oxidation-reduction

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Enthalpy and standard heats of reaction

Hess’s law

Spontaneous reactions and standard free energy change

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

Measurement of heat changes (calorimetry), heat capacity, specific heat

Heat transfer - conduction, convection, radiation

Bond dissociation energy as related to heats of formation

Spontaneous reactions and ΔG°

Coefficient of expansion

Heat of fusion, heat of vaporization

Phase diagram: pressure and temperature

Bioenergetics, thermochemistry and thermodynamics

Bioenergetics/thermodynamics

Free energy/Keq​

Concentration

Phosphorylation/ATP

Biological oxidation-reduction

Thermochemistry and thermodynamics

Endothermic and exothermic reactions

Enthalpy and standard heats of reaction

Hess’s law

Spontaneous reactions and standard free energy change

Phosphoryl group transfers and ATP

Biological oxidation-reduction

Thermodynamic system - state function

Zeroth law - concept of temperature

First law - conservation of energy in thermodynamic processes

PV diagram: work done = area under or enclosed by curve (PHY)

Second law - concept of entropy

Entropy as a measure of “disorder”

Relative entropy for gas, liquid, and crystal states

Measurement of heat changes (calorimetry), heat capacity, specific heat

Heat transfer - conduction, convection, radiation

Bond dissociation energy as related to heats of formation

Spontaneous reactions and ΔG°

Coefficient of expansion

Heat of fusion, heat of vaporization

Free energy/ $K_{e q}$

Spontaneous reactions and $Δ G °$

Free energy/ $K_{e q}$

Spontaneous reactions and $Δ G °$

Free energy/ $K_{e q}$

Spontaneous reactions and $Δ G °$

Free energy/ $K_{e q}$

Spontaneous reactions and $Δ G °$