Enthalpy (H): Total heat content of a system (kJ/mol)
System vs Surrounding: System = what we study, Surrounding = everything else
Activation Energy (Eₐ): Energy needed to start a reaction
Transition State: High-energy intermediate during reaction
Aerobic Respiration: With oxygen, produces CO₂ + H₂O + energy
Anaerobic Respiration: Without oxygen, produces ethanol/CO₂ + energy
Catalyst: Speeds up reaction by lowering activation energy
Bond Breaking: Endothermic (requires energy)
Bond Forming: Exothermic (releases energy)
Introduction to Chemical Energetics
Energy exists in different forms which are often interconvertible. In chemical energetics we are mainly concerned with two forms of energy:
1. Chemical Energy
This energy is stored in a molecule in which atoms are bonded to each other.
2. Heat Energy
This form of energy is released when a bond is formed and absorbed when it is broken.
In energetics we study the energy changes that take place during a chemical reaction. These changes are caused by the making and breaking of bonds during a reaction. In most of the reactions the weak bonds of reactants break while in products new strong bonds are formed.
Since energy is needed to break a bond while energy is evolved when a bond is formed, such reactions take place always with the evolution of heat. If a reaction is accompanied with the evolution of heat it is called an exothermic reaction and if heat is absorbed during a reaction it is called an endothermic reaction.
In energetics we not only encounter heat which comes out of a chemical reaction but also another quantity which is called enthalpy.
Enthalpy (H) or heat content, is defined as the total amount of thermal energy stored in a compound. The unit of its measurement is kJ mol⁻¹.
When energy is absorbed during a reaction, the total enthalpy of the system increases. When energy is evolved during a reaction, the total enthalpy of the system decreases.
Key Concepts – Introduction
Chemical Energy: Energy stored in chemical bonds
Heat Energy: Energy transferred due to temperature difference
Bond breaking → Endothermic (absorbs energy)
Bond forming → Exothermic (releases energy)
Enthalpy (H): Total heat content of system
ΔH: Change in enthalpy during reaction
Exothermic: ΔH = negative (heat released)
Endothermic: ΔH = positive (heat absorbed)
Memory Tips
EXOthermic = EXit of heat (heat leaves the system)
ENdothermic = ENter of heat (heat enters the system)
ΔH sign convention: Think of the system’s perspective
System loses heat → ΔH negative (-)
System gains heat → ΔH positive (+)
Bond Energy: Breaking bonds = input energy, Forming bonds = output energy
5.1 System and Surrounding
Thomas Young was the first to use the word ‘energy’ in the field of physics in 1802.
In Chemistry, any physical or chemical change under study may also be called a system. The chemical reaction includes reactants, products, catalyst, solvent and anything else which is important to study this reaction. Everything else which does not fall in this system is called the surrounding.
Example
If you are boiling water in a beaker, the water molecules will be called a system while everything surrounding this like beaker, burner, etc. will be called the surrounding.
When energy is transferred from surrounding to the system, the change is called endothermic and it has a positive sign. When energy is transferred from system to surrounding, the change is called exothermic and it carries a negative sign.
Interesting Information!
Energy evolved during a chemical reaction is used in everyday life for cooking, heating, lighting, transportation and much more.
Energy Transfer Between System and Surrounding
Endothermic Process
Heat flows INTO system
ΔH = POSITIVE (+)
Surroundings get cooler
Example: Boiling water
Exothermic Process
Heat flows OUT of system
ΔH = NEGATIVE (-)
Surroundings get warmer
Example: Burning fuel
Key Notes – System & Surrounding
System: Part of universe being studied (reactants, products, container)
Surrounding: Everything else in the universe
Endothermic: Heat flows FROM surrounding TO system
Exothermic: Heat flows FROM system TO surrounding
ΔH positive (+) = Endothermic (system gains heat)
ΔH negative (-) = Exothermic (system loses heat)
Temperature of surroundings decreases in endothermic process
Temperature of surroundings increases in exothermic process
Exercise 5.1
Does boiling water in a beaker represent an endothermic or exothermic change? Which form of energy is being transferred in this system?
Answer:
Boiling water is an endothermic change because heat energy is being transferred FROM the surroundings (burner, air) TO the system (water molecules).
The water molecules absorb heat energy to overcome intermolecular forces and change from liquid to gas phase. This causes the temperature of the surroundings (burner, air around beaker) to decrease slightly.
Energy form: Heat energy is being transferred.
5.2 Enthalpy
The total amount of heat energy present in a molecule under standard conditions (0°C temperature and 760 mm pressure) is also called its heat content. Enthalpy is the measurement of energy in a thermodynamic system. The quantity of enthalpy is equal to the total heat content of a system.
Enthalpy of a system is represented by (H) while the change in enthalpy which a system undergoes is represented by ΔH. The total enthalpy (H) of a system cannot be measured directly. However, the change in enthalpy ΔH brought about in a system can be measured comparatively easily.
In Chemistry, the standard enthalpy of reaction (ΔH) is the enthalpy change when reactants in their standard states undergo reaction to produce products in their standard states. This quantity is called the standard enthalpy change or heat of reaction at constant pressure.
Example Reaction
2CO + O₂ → 2CO₂
ΔH° = -566.0 kJ
The reaction in this system is thus exothermic evolving 566 kJ of heat energy which is given to the surrounding.
Interesting Information!
Enthalpy is important because it tells us how much heat is present in a system. Heat is important because we can extract useful work from it.
How is enthalpy different from heat?
HEAT
Form of energy that flows
Flows from hot to cold body
Measured in joules
Transfer of thermal energy
Not essential part of system
Comes and goes
Causes change in enthalpy
ENTHALPY
Property of a system
Total heat content
Measured in kJ/mol
Depends on composition/structure
Essential part of system
Constant for given system
Changes when heat transfers
Heat is a form of energy that flows from hot body to a cold body because of a difference in temperature. We measure heat in joules. Heat is what we call the transfer of thermal energy. Contrary to this, enthalpy is an essential part of a system since it depends on the number of molecules present in that system, its chemical composition and its structure. Heat is not essential part of a system, it just comes and goes. When heat leaves or enters a system, it results in a change of enthalpy. At a constant pressure the enthalpy change is equal to heat evolved or absorbed.
Key Notes – Enthalpy
Enthalpy (H): Total heat content of system
ΔH: Change in enthalpy during reaction
Cannot measure absolute enthalpy (H)
Can measure enthalpy change (ΔH)
Standard conditions: 0°C, 760 mmHg pressure
At constant pressure: ΔH = heat evolved/absorbed
Enthalpy depends on: Number of molecules, chemical composition, structure
Heat is energy transfer; Enthalpy is energy content
Exercise 5.2
1. Can energy be transferred in a form other than heat during a chemical reaction?
2. Why is it not possible to calculate the enthalpy of a system?
Answer 1: Yes, energy can be transferred in forms other than heat during chemical reactions. Examples include:
Light energy: In chemiluminescence reactions (glow sticks)
Electrical energy: In electrochemical cells (batteries)
Sound energy: In explosive reactions
Mechanical energy: In some decomposition reactions
Answer 2: It is not possible to calculate the absolute enthalpy (H) of a system because:
Enthalpy is a relative quantity, not an absolute one
We can only measure changes in enthalpy (ΔH), not the total enthalpy
There’s no zero point for enthalpy scale (like temperature has absolute zero)
We can measure enthalpy changes during reactions, but not the absolute value of H
5.3 Exothermic and Endothermic Reactions
A physical or a chemical change is almost always accompanied with either absorption or evolution of heat. Heat, which is evolved or absorbed during a chemical reaction, is called the heat of that reaction.
Chemical reactions in which heat energy is evolved are called exothermic reactions while those in which heat energy is absorbed are called endothermic reactions. Heat, which is evolved during an exothermic reaction, goes to the surrounding and the container in which such a reaction is being carried out, gets hot. Conversely, in an endothermic reaction, the absorption of heat from the surrounding will decrease the temperature of the container.
Example 1: Exothermic Reaction
2H₂(g) + O₂(g) → 2H₂O(l)
571.6 kJ heat energy is evolved during this reaction.
If the energy evolved is shown separately it is expressed as ΔH = -571.6 kJ.
The same amount of energy will be absorbed when the reaction will move in the backward direction i.e. water will decompose to give hydrogen and oxygen back.
Example 2: Exothermic Reaction
C(s) + O₂(g) → CO₂(g)
It is also an exothermic reaction and 393.5 kJ heat energy is evolved during this reaction.
When this reaction is carried out in the reverse direction, the same amount of energy i.e. 393.5 kJ will be absorbed. The enthalpy change (ΔH) for this reaction is -393.5 kJ mol⁻¹.
Interesting Information!
Heat evolved or absorbed during a reaction is used in self-heating or self-cooling packs. These packs contain reactants that undergo an exothermic or an endothermic reaction providing high or low temperature.
Endothermic Reactions
Example 3: Endothermic Reaction
H₂(g) + I₂(g) → 2HI(g)
Hydrogen gas reacts with iodine only at high temperature and 53.08 kJ of heat energy is absorbed.
The enthalpy change for the reaction is ΔH = +53.08 kJ.
Example 4: Endothermic Reaction
N₂(g) + O₂(g) → 2NO(g)
Formation of NO in air due to lightning in the clouds.
The enthalpy change for the reaction is ΔH = +180.6 kJ.
Comparing Exothermic vs Endothermic
EXOTHERMIC
Heat is RELEASED
ΔH = NEGATIVE (-)
Products have LESS energy than reactants
Container feels HOT
Surroundings temperature INCREASES
Examples: Combustion, neutralization
Bond forming > Bond breaking
ENDOTHERMIC
Heat is ABSORBED
ΔH = POSITIVE (+)
Products have MORE energy than reactants
Container feels COLD
Surroundings temperature DECREASES
Examples: Photosynthesis, boiling
Bond breaking > Bond forming
Key Notes – Exo/Endothermic Reactions
Exothermic: ΔH negative (-), Heat released
Endothermic: ΔH positive (+), Heat absorbed
Exothermic: Products have lower energy than reactants
Endothermic: Products have higher energy than reactants
Reverse reaction has opposite ΔH sign
Energy balance: ΔH = Energy of bonds broken – Energy of bonds formed
Most spontaneous reactions are exothermic (but not all)
Living organisms use exothermic reactions for energy
Quick Identification Tips
Exothermic Reactions (feel hot):
Combustion (burning)
Neutralization (acid + base)
Respiration
Most oxidation reactions
Freezing (liquid to solid)
Endothermic Reactions (feel cold):
Photosynthesis
Electrolysis
Evaporation/boiling
Most decomposition reactions
Melting (solid to liquid)
5.4 How does a Reaction take place?
A reaction takes place when the reactant molecules collide with each other to give a transition state. Let us study the following hypothetical reaction:
A₂ + B₂ → 2AB
Before mixing, the molecules of reactants A and B are in a state of random motion separately colliding with each other and with the walls of container. Kinetic energies possessed by these molecules are not the same. Majority of these molecules possess average kinetic energy but a few possess more than average energy while yet others possess less than average kinetic energy. The molecules which possess more than average kinetic energy may also be called excited molecules.
When the two reactant molecules are mixed together, all these molecules start colliding with each other. The collisions which result by colliding molecules having average or less than average kinetic energies may not be able to produce any result. But when the two excited molecules from both the reactants collide with each other they may be able to produce the transition state.
Reaction Pathway: A₂ + B₂ → 2AB
Formation of Transition State
A₂ + B₂
Reactants
[A-A-B-B]⁺
Transition State
AB + AB
Products
The transition state is a high-energy, unstable intermediate
Energy Profile Diagrams
Exothermic Reaction Energy Profile
Reactants
Products
Transition State
Eₐ (Activation Energy)
-ΔH (Heat Released)
Products have lower energy than reactants → Energy released → Exothermic
The energy of the transition state is higher than that of reactants or products because the bonds between the reactant or product molecules are being cleaved progressively. The energy absorbed by the reactant or product molecules in order to be converted into the transition state is called the activation energy (Eₐ) of the reaction. The difference between the energy of reactant and that of the product comes out in the form of heat representing enthalpy (ΔH) of the reaction.
Endothermic Reaction Energy Profile
Reactants
Products
Transition State
Eₐ (Activation Energy)
+ΔH (Heat Absorbed)
Products have higher energy than reactants → Energy absorbed → Endothermic
Interesting Information!
Washing clothes at 60°C uses almost twice the energy as at 30°C wash. 90% of the energy used by the traditional electric bulb is wasted in producing heat.
Role of Catalysts
An addition of the catalyst in a reaction increases the rate of reaction because it changes the path adopted by the reactants whereby the activation energy value of the reaction is substantially decreased. As a result, more reactants are now able to be converted into product molecules and hence the rate of reaction will increase.
Effect of Catalyst on Activation Energy
Reactants
Products
Eₐ (without catalyst)
Eₐ (with catalyst)
Catalyst provides alternative pathway with lower activation energy
A catalyst is thus, defined as a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. For example:
Ni acts as a catalyst in the hydrogenation of oil to give banaspati ghee
Platinum acts as a catalyst in the production of H₂SO₄
Chlorine acts as a catalyst promoting the breakdown of ozone
Key Notes – Reaction Mechanism
Collision Theory: Molecules must collide to react
Activation Energy (Eₐ): Minimum energy needed for reaction
Transition State: High-energy intermediate state
Only excited molecules (with E ≥ Eₐ) can form transition state
Exothermic: Products lower energy than reactants
Endothermic: Products higher energy than reactants
Catalyst: Lowers Eₐ, provides alternative pathway
Catalyst not consumed in reaction (regenerated)
Catalyst increases rate but not ΔH or equilibrium position
Exercise 5.4
1. Are energy diagrams useful?
2. Draw an energy profile diagram for a hypothetical reaction which does not evolve or absorb heat.
Answer 1: Yes, energy diagrams (energy profile diagrams) are extremely useful because:
They visually represent energy changes during reactions
They show activation energy (Eₐ) needed to start reaction
They indicate whether reaction is exothermic or endothermic
They show transition state (highest energy point)
They help compare catalyzed vs uncatalyzed reactions
They illustrate reaction pathways and intermediates
Answer 2: A reaction that doesn’t evolve or absorb heat has ΔH = 0. The energy profile diagram would show:
Energy Profile for ΔH = 0 Reaction
Reactants
Products
Transition State
Eₐ
Reactants and products at same energy level → ΔH = 0 → No heat evolved/absorbed
In such reactions, the energy of reactants equals the energy of products, so no net heat change occurs. The activation energy (Eₐ) is still needed to reach the transition state, but once products form, they have the same energy as the original reactants.
5.5 Aerobic and Anaerobic Respiration
The process of respiration in human beings is a continuous process. During this process, we breathe in oxygen and breathe out carbon dioxide. Respiration also carries complex chemical reactions inside the human body. This process that occurs in the presence of oxygen is called aerobic respiration.
Aerobic Respiration
Aerobic respiration is an exothermic process and involves the following reactions:
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy
(Glucose) (Oxygen) (Carbon dioxide) (Water)
ΔH = -2880 kJ/mol (approximately)
Stages of Aerobic Respiration
GLYCOLYSIS
Occurs in cytoplasm
Glucose → 2 Pyruvate
Catalysed by enzymes
Net production: 2 ATP
Doesn’t require oxygen
MITOCHONDRIA
Pyruvate → CO₂ + H₂O
Requires oxygen
Produces 36-38 ATP
Krebs cycle + ETC
Most ATP produced here
During glycolysis one molecule of glucose is split into two molecules of pyruvate. This process involves a series of reactions catalysed by enzymes, with a net production of 2 ATP (Adenosine Triphosphate). When cells of our body require energy for performing the metabolic activities, they use this ATP and break it down to get the required energy. The food we eat undergoes digestion in our body and the digested food molecules that are absorbed by the cells undergo oxidation to produce energy.
Anaerobic Respiration
In certain organisms like bacteria and algae respiration occurs in the absence of oxygen and it is called anaerobic respiration. This is also an exothermic process and during this process glucose is converted into carbon dioxide and ethanol with the evolution of energy.
C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂ + Energy
(Glucose) (Ethanol) (Carbon dioxide)
ΔH = -118 kJ/mol (approximately, much less than aerobic)
Energy Storage in Body
Lipids are a group of organic compounds which include fats, waxes, sterols, etc. Lipids serve as an energy reserve within our body. About half of the fuel our body needs comes from lipids. If you eat more food than you need in a day, the excess food is stored as lipids in adipose cells. In between meals and during exercise our body relies on this resource to provide energy.
Key Notes – Respiration
Aerobic Respiration: With oxygen, produces CO₂ + H₂O + energy
Anaerobic Respiration: Without oxygen, produces ethanol/CO₂ + energy
Both processes are exothermic (release energy)
Aerobic produces MUCH more energy (36-38 ATP vs 2 ATP)
ATP: Energy currency of cells
Glycolysis: First step in both types, occurs in cytoplasm
Lipids: Long-term energy storage (fats)
Glycogen: Short-term energy storage (in liver/muscles)
Respiration = Controlled “burning” of food for energy
Comparison Table: Aerobic vs Anaerobic
Feature
Aerobic Respiration
Anaerobic Respiration
Oxygen requirement
Requires oxygen
No oxygen required
Energy yield
High (36-38 ATP/glucose)
Low (2 ATP/glucose)
End products
CO₂ + H₂O
Ethanol + CO₂ or Lactic acid
Efficiency
High
Low
Examples
Human cells, animals
Yeast, bacteria, muscle cells during exercise
Complete Question Bank with Solutions
Multiple Choice Questions
1. Which of the following processes is endothermic?
(a) Burning of coal
(b) Respiration in humans
(c) Boiling of water
(d) Neutralization reaction
Answer: (c) Boiling of water
Explanation: Boiling water is endothermic because it absorbs heat from surroundings to convert liquid water to water vapor. The other processes are exothermic: burning coal releases heat, respiration releases energy, and neutralization reactions release heat when acid reacts with base.
2. In an exothermic reaction, the enthalpy change (ΔH) is:
(a) Zero
(b) Negative
(c) Positive
(d) Cannot be determined
Answer: (b) Negative
Explanation: In exothermic reactions, heat is released from the system to surroundings, so the system loses heat energy. By convention, when the system loses energy, ΔH is negative. For endothermic reactions (heat absorbed), ΔH is positive.
3. The energy required to start a chemical reaction is called:
(a) Enthalpy
(b) Activation energy
(c) Bond energy
(d) Kinetic energy
Answer: (b) Activation energy
Explanation: Activation energy (Eₐ) is the minimum energy that reactant molecules must possess to undergo a chemical reaction and form products. It’s the energy barrier that must be overcome for a reaction to occur.
4. Which of the following statements about catalysts is TRUE?
(a) They are consumed in the reaction
(b) They change the enthalpy of reaction
(c) They provide an alternative pathway with lower activation energy
(d) Both (a) and (b)
Answer: (c) They provide an alternative pathway with lower activation energy
Explanation: Catalysts work by providing an alternative reaction pathway with lower activation energy. They are not consumed in the reaction (regenerated at the end) and do not change the enthalpy (ΔH) of the reaction – they only change the rate, not the thermodynamics.
Questions for Short Answers
2(i) What is the difference between enthalpy and enthalpy change?
Answer:
Enthalpy (H) is the total heat content of a system at constant pressure. It’s an extensive property that depends on the amount of substance.
Enthalpy change (ΔH) is the difference in enthalpy between products and reactants in a chemical reaction. It represents the heat evolved or absorbed during the reaction at constant pressure.
Key difference: We cannot measure absolute enthalpy (H), but we can measure enthalpy changes (ΔH) during reactions.
2(ii) Why is breaking of a bond an endothermic process?
Answer: Bond breaking is endothermic because it requires energy input to overcome the attractive forces between atoms. When a bond breaks:
Energy must be supplied to separate the atoms
The atoms move from a lower energy state (bonded) to a higher energy state (separated)
This energy comes from the surroundings, so the process absorbs heat
Energy absorbed = bond dissociation energy
Example: Breaking H-H bond in H₂ molecule requires 436 kJ/mol of energy input.
2(iii) Depict the transition state for the following reaction: H₂ + Cl₂ → 2HCl
Answer: The transition state for H₂ + Cl₂ → 2HCl can be represented as:
[H-H-Cl-Cl]⁺ or [H···H···Cl···Cl]⁺
Explanation: In the transition state:
The H-H bond is partially broken
The Cl-Cl bond is partially broken
New H-Cl bonds are partially formed
All four atoms are weakly connected in a high-energy arrangement
The transition state exists for a very short time (femtoseconds)
It can collapse back to reactants or proceed to form products
2(v) What is the role of glycogen in our body?
Answer: Glycogen plays a crucial role as the short-term energy storage molecule in our body:
Storage form of glucose: Glycogen is a polysaccharide made of many glucose units
Storage sites: Primarily stored in liver and muscles
Liver glycogen: Maintains blood glucose levels between meals
Muscle glycogen: Provides immediate energy for muscle contraction during exercise
Quick energy release: Can be rapidly broken down to glucose when needed
Regulation: Insulin promotes glycogen synthesis after meals; glucagon promotes glycogen breakdown when blood sugar is low
Glycogen is like the body’s “quick-access” energy reserve, while fats are the long-term energy storage.
Constructed Response Questions
3(ii) Explain why the reaction between atmospheric gases oxygen and nitrogen does not take place under normal conditions? But in the presence of lightning these gases react to give NO. The reaction stops as soon as the lightning stops.
Answer:
The reaction N₂(g) + O₂(g) → 2NO(g) has a very high activation energy under normal conditions due to:
Strong triple bond in N₂: The N≡N bond has very high bond energy (945 kJ/mol), making it difficult to break
Strong double bond in O₂: The O=O bond also has high bond energy (498 kJ/mol)
Endothermic nature: The reaction is highly endothermic (ΔH = +180.6 kJ/mol), requiring substantial energy input
Why lightning enables the reaction:
High temperature: Lightning creates extremely high temperatures (≈30,000°C)
Energy source
Provides activation energy: The intense heat provides the necessary energy to overcome the high activation barrier
Excited molecules: High temperature increases kinetic energy, creating more molecules with energy ≥ Eₐ
Why reaction stops when lightning stops:
Temperature drops back to normal atmospheric levels
Insufficient energy to overcome activation barrier
Collisions between N₂ and O₂ molecules no longer have enough energy
Without continuous energy input, reaction cannot proceed
N₂(g) + O₂(g) → 2NO(g) ΔH = +180.6 kJ/mol
This explains both why nitrogen fixation by lightning occurs and why it’s limited to lightning events in nature.
Descriptive Questions
4(ii) Explain the difference between the terms heat and enthalpy.
Answer: Heat and enthalpy are related but distinct concepts in thermodynamics:
HEAT (q)
Form of energy transfer
Flows due to temperature difference
Measured in joules (J)
Path-dependent quantity
Not a property of system
Transient – comes and goes
Can do work
Example: Flame heating water
ENTHALPY (H)
Property of a system
Total heat content at constant pressure
Measured in kJ/mol
State function (path-independent)
Extensive property of system
Constant for given conditions
Cannot do work directly
Example: Energy stored in chemical bonds
Key Differences:
Nature: Heat is energy in transit; Enthalpy is stored energy
Measurement: We measure heat transfer; We measure enthalpy changes
Dependence: Heat depends on process path; Enthalpy depends only on initial and final states
Existence: Heat exists only during transfer; Enthalpy exists as long as the system exists
Relationship: At constant pressure, the enthalpy change (ΔH) equals the heat exchanged (qₚ): ΔH = qₚ
4(iii) Explain why formation of a bond is always an exothermic process.
Answer: Bond formation is always exothermic due to fundamental principles of energy and stability:
Energy Release Principle: When atoms come together to form a bond, they move from a higher energy state (separated atoms) to a lower energy state (bonded atoms). The excess energy is released as heat.
Stability Increase: Bonded atoms are more stable than separate atoms. The system lowers its potential energy by forming bonds, and this energy difference is released.
Electrostatic Attraction: As oppositely charged particles (nuclei and electrons) come closer, potential energy decreases, releasing energy.
Quantitative Explanation:
Separate atoms: High potential energy (unstable)
Bonded atoms: Low potential energy (stable)
Energy released = Bond dissociation energy
A + B → A-B + Energy
(Separate atoms) → (Bonded atoms) + Heat
Examples:
H + H → H₂ + 436 kJ/mol (H-H bond formation releases 436 kJ/mol)
Na⁺ + Cl⁻ → NaCl + 788 kJ/mol (Ionic bond formation in NaCl)
C + O₂ → CO₂ + 393.5 kJ/mol (Overall reaction involves bond formation)
Important Note: While individual bond formation is always exothermic, overall reactions may be endothermic if more energy is required to break existing bonds than is released by forming new bonds.
4(v) Explain the following terms: Activation energy, Transition state, Aerobic respiration
Answer:
Activation Energy (Eₐ)
Definition: The minimum amount of energy that reactant molecules must possess to undergo a chemical reaction and form products.