Chemistry Hougang: Thermodynamics Explained

Did you know? Some chemical changes release heat, yet that does not automatically tell us whether the process is thermodynamically favourable. Other changes absorb heat while proceeding naturally under particular conditions. A battery can deliver electrical work, but not all the energy associated with its chemical reaction becomes useful electricity. These are questions that thermodynamics helps us understand.

This Chemistry Hougang guide explains chemical thermodynamics, enthalpy, entropy, Gibbs free energy, reaction spontaneity, equilibrium, and the importance of energy accounting in real technology. It is written for Secondary learners building foundations, JC H1 and H2 Chemistry students studying energetic changes, and parents searching for a Chemistry tutor in Hougang who teaches why the equations make sense.

I want students to recognise four separate ideas: energy is conserved, energy can be transferred as heat or work, chemical changes have a direction favoured under specified conditions, and an energetically favoured process may still be slow. Those statements are connected, but none should be mistaken for another.

The wider world needs these distinctions. They help us evaluate batteries, industrial reactors, solar technologies, desalination, climate proposals and the enormous energy exchanges supporting living systems.

What is thermodynamics?

Thermodynamics studies energy, heat, work and the relationships among the states and changes of physical and chemical systems. Chemical thermodynamics uses properties such as enthalpy, entropy and Gibbs free energy to describe reactions and phase changes.

The OpenStax Chemistry introduction to thermodynamics explains how thermodynamics helps predict whether a process is favoured under specified conditions, while not automatically describing how quickly it happens.

I find this contrast important. In Chemistry, thermodynamics addresses state-related driving forces; kinetics addresses rates and pathways.

Imagine a reaction with products thermodynamically favoured under the stated conditions. If the activation barrier is very high, the reaction may still occur extremely slowly without a suitable pathway.

A student can therefore be correct in saying that a process is spontaneous in the thermodynamic sense and also correct in saying that its rate is negligible over the time being considered.

That is not a contradiction. It is the first major conceptual boundary students must learn.

What does spontaneous actually mean?

In everyday language, “spontaneous” may sound like “sudden”, “explosive” or “happens immediately”. In thermodynamics it has a more precise meaning.

A thermodynamically spontaneous process is favoured in a stated direction under particular conditions. That does not mean it must proceed rapidly or that a visible event begins immediately.

The OpenStax thermodynamics summary emphasises that some spontaneous changes are slow and that spontaneity must be distinguished from the rate of change.

For example, a surface can undergo thermodynamically favourable corrosion while degradation happens slowly. A mixture may have a more stable state but remain trapped by kinetic barriers for a long period.

I would ask a student to replace the vague phrase “this reaction happens by itself” with a more careful sentence: the process is thermodynamically favoured in the stated direction under these conditions, but the rate requires a separate explanation.

This one correction can improve answers about reaction energetics, catalysts, crystallisation and electrochemistry.

The first law: energy does not appear from nowhere

The first law of thermodynamics expresses conservation of energy. A system can exchange energy with its surroundings, including as heat and work, but energy is accounted for rather than created or destroyed in an ordinary process.

The sign convention for heat and work must be specified. A common Chemistry convention writes the change in internal energy as:

ΔU = q + w, where q represents heat transferred into the system and w is work done on the system.

Other engineering conventions may choose different signs for work, so the meaning of each symbol matters more than memorising a symbol without its definition.

The OpenStax first-law chapter introduces heat, work and internal-energy changes using the Chemistry sign convention.

I tell students that an exothermic reaction is not “making energy”. Energy is transferred from the chemical system to its surroundings in the relevant process.

A reaction can change which forms and locations hold energy while the total accounting remains consistent.

What is the difference between heat and temperature?

Temperature is a state variable associated with the thermal condition of matter. Heat is energy transferred because of a temperature difference.

A hot, small object does not necessarily contain more total thermal energy than a much larger, cooler object. Amount of matter, heat capacity and the temperature reference matter.

Students sometimes treat Celsius degrees, joules and kilojoules as if they were interchangeable energy units. They are not.

Temperature differences may be expressed in °C or K with equal numerical increments, but absolute thermodynamic temperatures used in formulas such as ΔG = ΔH − TΔS must be in kelvin.

The difference between a temperature reading and an energy amount is also vital for industrial heating, refrigeration and climate systems.

I ask learners to write the correct units beside every number before beginning a thermodynamics problem. A correct unit often reveals whether the chosen calculation has the right physical meaning.

Worked example 1: heating water

Imagine 2.0 kg of water heated from 20°C to 35°C, with a simplified constant specific heat capacity of 4.2 kJ kg⁻¹ K⁻¹ and no phase change.

The temperature rise is:

ΔT = 35 − 20 = 15 K.

Using q = mcΔT:

q = 2.0 kg × 4.2 kJ kg⁻¹ K⁻¹ × 15 K = 126 kJ.

The units kg and K cancel, leaving kilojoules of energy.

This calculation estimates the sensible heat gained by the water. It does not determine the electrical energy drawn from a real heater if there are losses to the environment or additional objects being heated.

A child who simply says “the answer is 126 degrees” has confused energy with temperature. A child who gets 126 kJ but claims no external energy is required has made a different conceptual mistake.

The purpose of this example is to connect equation, unit and physical meaning.

Enthalpy: why chemists use ΔH

Enthalpy, H, is a thermodynamic state function defined as H = U + pV, where U is internal energy, p pressure and V volume.

Under particular conditions, the change in enthalpy corresponds to heat exchanged at constant pressure when only pressure–volume work is relevant. This makes ΔH useful for describing many reaction and phase-change energy changes.

In school Chemistry, we often use negative ΔH for an exothermic process and positive ΔH for an endothermic process under the familiar sign convention.

An exothermic reaction transfers heat to the surroundings under the relevant conditions. An endothermic process requires heat transfer into the system under the corresponding description.

But ΔH does not tell us everything about spontaneity. It does not, by itself, measure reaction speed or final product yield.

The OpenStax thermochemistry discussion explains reaction enthalpies and why they are useful for energy balances.

I want learners to understand that the sign of ΔH is a property of a defined change between states, not a moral judgment about whether a reaction is good.

Why melting and evaporation are thermodynamic changes

Melting and evaporation involve changes of physical state. The chemical formula of the substance may stay the same while the arrangements and interactions among particles change.

A solid can melt at a suitable temperature and pressure, or a liquid can evaporate without needing to become a different chemical compound.

These changes involve energy and entropy. Their thermodynamic favourability depends on conditions.

Students sometimes label an endothermic melting process as “impossible without a chemical reaction”. That is incorrect. An endothermic phase change can be thermodynamically favoured under suitable temperatures.

Our crystallisation guide shows how the reverse formation of ordered solids depends on solubility and conditions. Our refrigeration guide explains how designed phase changes can be used in energy-transfer systems.

This is a powerful link between particle models, matter’s physical states and modern technology.

Entropy: more useful than “disorder”

Entropy, S, is a thermodynamic state function related to how energy and matter are distributed among accessible microscopic states. It is frequently introduced as “disorder”, but that word can be misleading if used as a universal definition.

The OpenStax thermodynamics summary connects entropy to microscopic arrangements and dispersal of matter and energy.

An increased number of accessible microstates can help explain why entropy may increase during certain changes, such as mixing or vaporisation under suitable conditions.

But entropy is not simply a score of how untidy a room looks. A perfectly arranged bookshelf and a messy desk are not a scientifically reliable way to calculate thermodynamic entropy.

I would rather a Secondary learner know that entropy concerns the distribution of matter and energy, then add rigorous equations as mathematical preparation improves.

At university, statistical thermodynamics connects entropy to microstate counts and probabilities with greater precision.

The second law: remember the whole system

The second law of thermodynamics states that, for a spontaneous process in an isolated total system, total entropy increases; at equilibrium there is no net entropy production in the idealised equilibrium description.

This is often summarised as ΔS_universe > 0 for a spontaneous process under the relevant framework.

The word universe in this equation refers to the system and its surroundings as a complete accounting boundary—not necessarily every galaxy as something the student must measure directly.

The entropy of the particular system can decrease while the entropy change of the surroundings is large enough to make the total change positive.

For example, an ordered solid may form from a liquid under suitable conditions, reducing the system’s entropy. The process can still be spontaneous if the total thermodynamic accounting favours it.

This is why the statement “a spontaneous process must always make the system more disordered” is false.

Good Chemistry requires carefully defining the system and surroundings before interpreting the second law.

Why Gibbs free energy is such a useful quantity

At constant temperature and pressure, the Gibbs free energy change, ΔG, helps describe the direction of a chemical or physical change.

The familiar expression is:

ΔG = ΔH − TΔS.

Here T is absolute temperature in kelvin, and the energy units for ΔH and TΔS must match.

Under the relevant conditions, ΔG < 0 indicates a forward change that is thermodynamically favoured; ΔG > 0 indicates that the forward direction is not favoured; and ΔG = 0 characterises an equilibrium condition for the specified transformation.

The OpenStax explanation of free energy develops these relationships and connects ΔG to the entropy change of the system plus surroundings.

Students should understand the conditions and direction represented. A positive ΔG for the forward change does not mean the reverse change is also thermodynamically disfavoured.

Nor does a negative ΔG mean a reaction proceeds instantly or produces every possible product without competition.

Worked example 2: enthalpy and entropy compete

Consider a fictional chemical change with:

ΔH = +25.0 kJ mol⁻¹

and

ΔS = +100 J mol⁻¹ K⁻¹.

First convert entropy units:

100 J mol⁻¹ K⁻¹ = 0.100 kJ mol⁻¹ K⁻¹.

At T = 298 K, calculate:

ΔG = 25.0 − (298 × 0.100) = 25.0 − 29.8 = −4.8 kJ mol⁻¹.

So the forward change is thermodynamically favoured under the stated constant-temperature-and-pressure model.

The process is endothermic because ΔH is positive, yet its ΔG is negative because the positive entropy contribution is sufficiently influential at this temperature.

This is a direct correction to the misconception “every endothermic process must be nonspontaneous”.

I want students to pay attention to the J versus kJ conversion. Omitting the factor of 1,000 would create an enormous error.

Worked example 3: find the crossover temperature

Using the same fictional reaction, solve for the temperature where ΔG = 0, assuming ΔH and ΔS are approximately constant over the range:

0 = ΔH − TΔS.

Therefore:

T = ΔH / ΔS = 25.0 / 0.100 = 250 K.

Above this crossover temperature in the simplified model, ΔG is negative; below it, ΔG is positive.

At 200 K, for example:

ΔG = 25.0 − (200 × 0.100) = +5.0 kJ mol⁻¹.

The forward process is not thermodynamically favoured under those stated conditions.

This does not mean the process suddenly becomes infinitely fast at 251 K. We have not calculated any kinetic barrier or rate constant.

The crossover is a thermodynamic sign change, not a prediction of reaction speed.

At advanced levels, ΔH and ΔS may vary with temperature, and phase changes or other complications can require a more complete model.

The four sign combinations make a useful map

At the usual school approximation of ΔG = ΔH − TΔS, with T positive:

ΔH negative, ΔS positive: the forward process is thermodynamically favoured for all positive temperatures within the model’s validity.
ΔH positive, ΔS negative: the forward process is not thermodynamically favoured for positive temperatures within the same model.
ΔH positive, ΔS positive: favourability may depend on temperature, often favoured at sufficiently high T.
ΔH negative, ΔS negative: favourability may depend on temperature, often favoured at sufficiently low T.

The OpenStax free-energy chapter explains these categories with examples.

I teach students not to memorise this table without understanding the negative sign in −TΔS. The sign logic becomes straightforward after units and values are written consistently.

The table is not a universal guarantee across all temperatures, because material phases and thermodynamic quantities can change as conditions vary.

It is a model with useful domain conditions.

Gibbs free energy and reaction equilibrium

Under appropriate thermodynamic conventions, the standard free-energy change is linked to an equilibrium constant:

ΔG° = −RT ln K.

The actual reaction free-energy change away from standard-state conditions is related to the reaction quotient Q:

ΔG = ΔG° + RT ln Q.

R is the gas constant, T is absolute temperature and ln is the natural logarithm. K and Q are defined using dimensionless activities in a rigorous treatment.

When a system is at equilibrium, Q = K and the reaction’s ΔG is zero.

Our chemical equilibrium guide explains why equilibrium composition and reaction rate are different questions.

An important distinction is that ΔG° does not have to be zero when a system is at equilibrium. It depends on the standard-state reference; the actual ΔG is zero at equilibrium under the appropriate conditions.

That difference often causes confusion in JC Chemistry problems.

Worked example 4: equilibrium constant and ΔG°

For a fictional reaction at 298 K, suppose the dimensionless equilibrium constant is K = 100.

Use R = 8.314 J mol⁻¹ K⁻¹:

ΔG° = −(8.314)(298) ln(100).

Since ln(100) ≈ 4.605, this gives:

ΔG° ≈ −11,400 J mol⁻¹ = −11.4 kJ mol⁻¹.

The negative standard free-energy change is consistent with K greater than one for the reaction as written.

But it does not mean that every mixture must currently undergo a net forward reaction. The actual direction depends on the present reaction quotient Q.

If a particular mixture already satisfies Q = K, the system is at equilibrium even though ΔG° is negative.

I like this example because it moves students beyond the shortcut that “negative always means products keep forming forever”.

Why catalysts cannot change the equilibrium constant

A catalyst provides a different kinetic pathway and can allow a reaction to proceed more rapidly.

Under the same temperature and underlying thermodynamic conditions, it does not change the equilibrium constant or the thermodynamic free-energy difference between the defined initial and final states.

Students often confuse “the catalyst increased daily production” with “the catalyst changed K”. Real industrial throughput can improve when a reaction reaches equilibrium more quickly or works effectively under chosen conditions, even though K remains fixed by thermodynamics at that temperature.

Our catalysis guide examines activation energy and rate.

This is one of the most valuable connections in Chemistry: thermodynamics tells us the destination favoured by the specified conditions; kinetics helps determine whether and how quickly the system can get there.

A tutor should ask learners to explain both sentences without mixing the terms.

Hess’s law: why the route should not change ΔH

Hess’s law follows from enthalpy being a state function. If the same initial and final thermodynamic states are connected by different imagined reaction pathways, the total enthalpy change is the same.

Imagine a fictional transformation A → C that can be represented through an intermediate B:

A → B: ΔH = −50 kJ mol⁻¹ on the defined common reaction basis.

B → C: ΔH = +20 kJ mol⁻¹ on that same basis.

The total change is:

ΔH(A → C) = −50 + 20 = −30 kJ mol⁻¹.

A direct A → C transformation would have the same ΔH between the same initial and final states.

This does not mean the reaction rates or experimental feasibility of the two routes are identical. The intermediates and activation barriers can differ greatly.

Hess’s law is powerful because it allows students to combine thermochemical information while recognising what energy accounting can and cannot reveal.

Why calorimetry needs careful measurement

Calorimetry measures heat transfer under defined experimental arrangements. In introductory experiments, students may use temperature changes and heat capacities to estimate energy transferred.

But the outcome depends on insulation, apparatus heat capacity, sample composition and assumptions about what exchanges heat.

A student might accurately record a 4°C temperature rise but then incorrectly treat 4°C as the enthalpy change of a reaction. The measured temperature change must be converted into energy using relevant quantities, and a reaction enthalpy may require dividing by amount reacted.

The result should also use appropriate sign conventions for the chemical system and surroundings.

Our chemical analysis guide explains why a measurement is only as useful as the method and interpretation.

A calorimetry number that looks precise can still be biased if the system loses heat or a relevant object was omitted from the calculation.

Thermodynamics in batteries

A battery can convert chemical free energy into electrical work under appropriate conditions. The useful voltage is connected to electrochemical driving forces, but actual performance also involves resistance, transport, heat and other losses.

A negative Gibbs free-energy change indicates a driving force for a particular reaction under given conditions; it does not guarantee that a device extracts every theoretically available joule as electricity.

The OpenStax free-energy section relates ΔG to the theoretical maximum useful non-expansion work under the appropriate reversible idealisation.

Our Lithium-Ion Batteries guide explains electrons, ions, charging and degradation.

The engineering lesson is that a theoretical thermodynamic limit and a real product rating are different quantities. A professionally useful device must also be safe, reliable and suited to the application.

Students should not attempt battery disassembly or homemade cell experiments based on a thermodynamics calculation.

Thermodynamics in hydrogen and electrolysis

Splitting water to make hydrogen requires an external energy input under the relevant conditions. Hydrogen can then serve as an energy carrier.

The Hydrogen Fuel and Electrolysis guide explains the simplified reaction:

2H₂O → 2H₂ + O₂.

A balanced equation conserves atoms. Thermodynamics helps describe the energy requirements and driving forces. Kinetics and engineering determine additional real losses and useful production rates.

It would be misleading to say that a catalyst makes an electrolysis process “free”. Catalysts can improve rates and reduce kinetic overpotentials but cannot eliminate thermodynamic requirements or create energy.

This is why energy technologies must be judged using full system boundaries—electrical inputs, chemical-energy output, conversion stages, storage and distribution.

A student’s correct half-equation is the beginning of a system analysis, not the final policy conclusion.

Thermodynamics and refrigeration

A refrigerator moves heat from a colder region to a warmer surrounding environment with an appropriate external work input.

That can seem counterintuitive because heat naturally flows from hot to cold when left to ordinary unassisted conduction. A refrigerator uses a designed cycle and energy input to move heat in the opposite direction.

The Refrigeration and Heat Pumps guide explains refrigerant phase changes and energy transfer.

A heat pump’s coefficient of performance can exceed one because it is moving existing heat as well as using supplied energy. It is not creating energy from nothing.

This distinction is essential when reading claims about efficient heating and cooling. An efficiency-like number above 100% may describe a heat-transfer ratio, not an energy-generation efficiency.

The correct quantity must be named and understood before comparison.

Thermodynamics of dissolving and crystallising

A substance can dissolve in water even when the dissolution absorbs heat under some conditions. Conversely, crystals can form spontaneously from an appropriate solution under other conditions.

This becomes understandable once we consider enthalpy and entropy rather than looking only at heat.

For a specified system, the interplay of solute–solute, solvent–solvent and solute–solvent interactions, temperature and composition affects dissolution and equilibrium.

Our crystallisation guide explains saturation, supersaturation and solid-form control.

A student should not infer that an endothermic dissolution process must be impossible. Nor should they assume that a thermodynamically favoured solid necessarily crystallises instantly: nucleation can require a kinetic step.

That is a perfect everyday demonstration of why both thermodynamic and kinetic reasoning are needed.

Living systems: order can grow while entropy increases overall

Living organisms maintain complex internal structures through continuous energy and material flows. Photosynthesis, respiration and metabolism involve transformations that connect thermodynamic driving forces with enzymes and biological regulation.

The existence of ordered cells does not contradict the second law. Organisms are not isolated systems with no exchange of energy or matter.

Maintaining and building local organisation can be accompanied by entropy production and energy dispersal elsewhere in the surroundings.

Our Photosynthesis guide explains how light energy helps sustain organic chemical transformations, while our food Chemistry guide develops fermentation and enzyme examples.

Students should avoid claiming that a plant “creates energy from sunlight”. It converts incoming radiation energy into forms useful to biological systems, with losses and transfers accounted for.

The broader point is that life operates within thermodynamics, not outside it.

Climate and energy decisions need more than one thermodynamic number

Reducing emissions, improving process efficiency and choosing energy technologies all benefit from careful energy and material accounting.

But a favourable reaction enthalpy does not prove that a technology is low-emissions. The full process may require raw-material production, transport, separation, cooling, maintenance and waste management.

Similarly, a process advertised as “energy efficient” must be compared using consistent service requirements and measurement boundaries.

Our carbon capture guide explains why a high percentage of captured CO₂ does not automatically equal an identical reduction in net facility emissions.

Our chemical engineering guide discusses energy and mass balances at industrial scale.

Thermodynamics supplies essential constraints. Engineering and environmental analysis determine how a useful system actually performs in the world.

That is the difference between understanding a reaction and making a responsible public recommendation.

What a Chemistry tutor should diagnose

Imagine a student says a reaction has negative ΔH, therefore its ΔG must also be negative. That student has conflated enthalpy and Gibbs free energy.

Another student uses ΔS = 100 J mol⁻¹ K⁻¹ directly beside ΔH = 25 kJ mol⁻¹ and forgets to convert units. That student may understand the concept but need stronger numerical discipline.

A third student correctly calculates ΔG < 0 and then claims the reaction will finish in two seconds. That student needs to distinguish thermodynamics from kinetics.

Those are three different learning gaps. A thoughtful Chemistry tutor for a Hougang student should diagnose each separately rather than assign the same set of worksheet repetitions.

I use one energy diagram, one numerical problem and one unfamiliar transfer question to see whether understanding has improved.

The best indicator is that the learner can explain a boundary without waiting for the teacher to prompt them.

Learning progression through Secondary and JC

Lower Secondary: energy and changes

Distinguish heat, temperature, physical change and chemical reaction. Learn the meaning of conserved energy and correct units.

Upper Secondary: exothermic and endothermic

Work with simple energy profiles, reaction energetics and appropriate calculations. Recognise that ΔH describes an energy change rather than reaction speed.

JC H1/H2: thermodynamics and quantitative analysis

Use ΔG = ΔH − TΔS, sign logic, Hess’s law and related equilibrium concepts at the level required by the applicable syllabus. Practise unit conversions and explanation of conditions.

Polytechnic, university and professional work

Study chemical potentials, nonideal mixtures, heat engines, process energy balances, statistical thermodynamics, electrochemistry and phase behaviour.

The scientific progression is continuous. Students build from basic observation toward quantitative models with explicit assumptions and evidence.

A four-week study plan without experimental hazards

Week 1 — Distinguish quantities. Write a short explanation of temperature, heat, internal energy and enthalpy. State the correct unit for each quantity in a typical school problem.

Week 2 — Calculate. Solve the water-heating example and the fictional ΔH = +25 kJ mol⁻¹, ΔS = +100 J mol⁻¹ K⁻¹ case at 298 K. Check your conversion to kJ.

Week 3 — Thermodynamics versus kinetics. Explain why a catalyst can change reaction speed without changing ΔG between the defined chemical states. Read the OpenStax free-energy and catalysis explanations.

Week 4 — World view. Compare a battery, a refrigerator and a crystallising solution. Explain what energy enters, what useful effect occurs, and which additional measurements would be needed for a real performance claim.

No heating of unknown substances, battery work or pressurised equipment is required for these exercises.

At the end, ask the student to answer: Can an endothermic process be spontaneous, and can a spontaneous process be slow? If they explain both correctly, the core ideas are becoming secure.

Common thermodynamics misconceptions

“Exothermic means spontaneous.” Enthalpy alone does not determine Gibbs free-energy sign.
“Spontaneous means fast.” Thermodynamic favourability and reaction kinetics are different.
“Entropy is just untidiness.” Entropy has a physical definition related to accessible states and energy or matter distribution.
“T in ΔG is Celsius.” Thermodynamic temperature must be expressed in kelvin.
“A catalyst can change the equilibrium constant at the same temperature.” A catalyst changes kinetic pathways, not the underlying equilibrium constant.
“A heat pump producing more heat than input electricity breaks energy conservation.” It transfers heat from an external source as well as using work input.

I would rather a student explain why each statement fails than memorise another chart without understanding the equations.

Self-check with model answers

If ΔH = +20 kJ/mol and ΔS = +80 J/mol/K at T = 300 K, what is ΔG?

Convert 80 J to 0.080 kJ/K/mol:

ΔG = 20 − 300(0.080) = −4 kJ/mol.

Is this reaction exothermic?

No. Its ΔH is positive, so it is endothermic under the stated convention.

Is its forward process thermodynamically favoured at 300 K?

Yes, because the calculated ΔG is negative under the stated model.

Does this result determine reaction speed?

No. It does not provide a rate constant or activation barrier.

What does ΔG = 0 represent at equilibrium?

For the defined transformation at the actual system conditions, it means no net free-energy driving force for reaction progress.

If a process has an ideal maximum useful electrical work of 100 units, must a real device deliver all 100?

No. Real devices have irreversible losses and other engineering limits.

Questions parents and students often ask

Why is thermodynamics so difficult in JC Chemistry?

Students must coordinate concepts, algebra, sign conventions, units and the boundary conditions of each model. A weakness in any one can make a simple-looking problem confusing.

Should a Secondary student learn Gibbs free energy early?

Begin with energy changes and careful units. The deeper equations become useful once the learner has sufficient Mathematics and syllabus foundations.

Does thermodynamics lead to university careers?

Yes. It supports Chemistry, chemical engineering, materials, energy systems, environmental science, biomedical research and many other fields.

What is the most useful revision habit?

After every calculation, state what the numerical answer means physically and which assumptions were made.

Sources and further reading

Continue the Chemistry Hougang learning library

A student who can distinguish ΔH, ΔG and reaction rate has built a strong foundation for many advanced topics. Families exploring support can arrange a Chemistry learning consultation with eduKate Singapore.

Chemistry Hougang | Thermodynamics Explained: Enthalpy, Entropy and Chemical Energy