Chemical Reactions 3D
Chemical Reactions 3D is a free interactive simulation from StudyFi Lab. It runs right in your browser, with no sign-up and nothing to install.
Watch atoms rearrange as bonds break and form
What you can explore
Formation of water
2 H₂ + O₂ → 2 H₂O
The H–H bonds in hydrogen molecules and the O=O double bond in oxygen break, and each oxygen atom forms two new O–H bonds. The new bonds are stronger than the broken ones, so a lot of energy is released – a mixture of hydrogen and oxygen explodes once ignited.
In everyday life: Fuel cells and some rocket engines use this reaction to get energy from hydrogen, and the only product is water.
Synthesis of ammonia
N₂ + 3 H₂ → 2 NH₃
The very strong N≡N triple bond in nitrogen and the H–H bonds in hydrogen break, and each nitrogen atom forms three N–H bonds. The triple bond is so hard to break that the reaction only runs at a useful rate with a catalyst; the Haber–Bosch process makes ammonia for the fertilisers that help feed a large part of humanity.
In everyday life: Almost all nitrogen fertilisers spread on fields are made from ammonia produced by this reaction.
Synthesis of hydrogen chloride
H₂ + Cl₂ → 2 HCl
Light or heat splits the Cl–Cl bond into chlorine atoms (radicals), which start a chain reaction: H–H and Cl–Cl bonds break and H–Cl bonds form. The reaction releases energy, and in strong light the mixture can even explode; dissolving HCl gas in water gives hydrochloric acid.
In everyday life: Hydrochloric acid is used to remove limescale and clean metals, and a dilute form of it is also in your gastric juice.
Formation of nitrogen monoxide
N₂ + O₂ → 2 NO
The N≡N triple bond in nitrogen and the O=O double bond in oxygen break, and nitrogen–oxygen bonds form in two NO molecules. Breaking these strong bonds takes more energy than the new bonds give back, so the reaction absorbs heat and only happens at very high temperatures – which is why the nitrogen and oxygen in air normally do not react.
In everyday life: Nitrogen monoxide forms in lightning and in hot car engines, and it is the first step towards the nitrogen oxides in exhaust fumes and smog.
Formation of nitrogen dioxide
2 NO + O₂ → 2 NO₂
The O=O double bond in oxygen breaks and each NO molecule binds one more oxygen atom through a new nitrogen–oxygen bond. The reaction runs on its own at room temperature and releases heat: colourless NO turns into reddish-brown, toxic NO₂.
In everyday life: The brownish haze over busy cities on smoggy days is partly NO₂ formed from exhaust gases, and it also contributes to acid rain.
Oxidation of sulfur dioxide
2 SO₂ + O₂ → 2 SO₃
The O=O double bond in oxygen breaks and each SO₂ molecule gains a third oxygen atom through a new sulfur–oxygen bond. The reaction releases heat and is reversible, so industry speeds it up with a vanadium(V) oxide catalyst; it is the key step of the contact process for making sulfuric acid.
In everyday life: Sulfuric acid made this way goes into fertilisers and car batteries, and the same oxidation of SO₂ in the air helps cause acid rain.
Combustion of carbon monoxide
2 CO + O₂ → 2 CO₂
The O=O double bond in oxygen breaks and each CO molecule takes up one oxygen atom; the carbon then holds two C=O double bonds in a linear CO₂ molecule. CO burns with a blue flame and releases a lot of heat – and the poisonous CO becomes much less dangerous CO₂.
In everyday life: The catalytic converter in a car's exhaust uses this reaction to remove poisonous carbon monoxide from the exhaust gases.
Formation of carbonic acid
CO₂ + H₂O → H₂CO₃
A water molecule attaches to the carbon atom of CO₂: one C=O double bond becomes a single bond, an O–H bond in water breaks and its hydrogen moves to an oxygen, so H₂CO₃ ends up with two O–H groups. It is an equilibrium lying far on the side of the reactants – most CO₂ in water simply stays dissolved and only a small part becomes carbonic acid, which makes the solution weakly acidic.
In everyday life: Sparkling drinks taste slightly sour because of carbonic acid, and when you open the bottle the dissolved CO₂ escapes as bubbles.
Decomposition of hydrogen peroxide
2 H₂O₂ → 2 H₂O + O₂
The weak O–O single bond and some O–H bonds in H₂O₂ break; the atoms regroup into water molecules with new O–H bonds and an oxygen molecule with an O=O double bond. On its own the reaction is slow, but a catalyst such as manganese dioxide or the enzyme catalase speeds it up enormously, and the released oxygen makes the mixture foam.
In everyday life: When you pour hydrogen peroxide on a cut, it fizzes because catalase in your blood breaks it down into water and oxygen.
Electrolysis of water
2 H₂O → 2 H₂ + O₂
Electric current supplies the energy needed to break the O–H bonds in water; the hydrogen atoms pair up through H–H bonds and the oxygen atoms through an O=O double bond. It is the reverse of the formation of water, so it needs at least as much energy as burning hydrogen releases; hydrogen forms at the cathode and oxygen at the anode in a 2 : 1 volume ratio.
In everyday life: Electrolysers powered by solar or wind electricity use this reaction to produce “green” hydrogen as a clean fuel.
Decomposition of ozone
2 O₃ → 3 O₂
In ozone, three oxygen atoms are held together by two bonds that are weaker than the O=O double bond; UV light breaks them, and the freed atoms rearrange into three O₂ molecules with O=O double bonds. Overall energy is released, and in the ozone layer the ozone molecules absorb the Sun's dangerous ultraviolet radiation as they are split.
In everyday life: The ozone layer high in the stratosphere protects you from sunburn and skin damage by absorbing UV light this way.
Decomposition of dinitrogen tetroxide
N₂O₄ → 2 NO₂
The relatively weak N–N bond in the middle of the N₂O₄ molecule breaks, splitting it into two NO₂ molecules; no new bonds form, so the reaction absorbs energy. It is a reversible equilibrium: warming shifts it towards brown NO₂, while cooling brings back colourless N₂O₄ – a classic demonstration of Le Chatelier's principle.
In everyday life: In the school lab, a sealed tube of this gas turns darker in hot water and paler in ice water; N₂O₄ is also used as an oxidiser in spacecraft engines.
Combustion of methane
CH₄ + 2 O₂ → CO₂ + 2 H₂O
All four C–H bonds in methane and the O=O double bonds in two oxygen molecules break; the carbon atom forms two C=O double bonds in CO₂ and the hydrogen atoms form O–H bonds in two water molecules. Together the new bonds are much stronger than the broken ones, so a lot of heat is released – which is why natural gas is such a good fuel.
In everyday life: The blue flame of a gas stove or boiler is methane from natural gas burning.
Combustion of ethane
2 C₂H₆ + 7 O₂ → 4 CO₂ + 6 H₂O
In each ethane molecule the C–C bond and six C–H bonds break, together with the O=O bonds of the oxygen molecules; the carbon atoms end up in CO₂ (C=O bonds) and the hydrogen atoms in water (O–H bonds). Two ethane molecules need seven oxygen molecules, and like every alkane combustion the reaction releases a lot of heat.
In everyday life: Ethane is the second most abundant component of natural gas after methane, so it burns alongside it in gas stoves and power plants.
Combustion of propane
C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O
The two C–C bonds and eight C–H bonds of propane break, together with the O=O bonds of five oxygen molecules; three CO₂ molecules form with C=O double bonds and four water molecules with O–H bonds. Propane is easily liquefied under moderate pressure, so a lot of energy fits into a small gas cylinder.
In everyday life: Camping stoves, gas barbecues and LPG cars burn propane (often mixed with butane).
Combustion of ethanol
C₂H₅OH + 3 O₂ → 2 CO₂ + 3 H₂O
The C–C, C–H, C–O and O–H bonds of ethanol break along with the O=O bonds of three oxygen molecules; the atoms recombine into two CO₂ molecules (C=O bonds) and three water molecules (O–H bonds). Because ethanol already contains one oxygen atom, it needs less oxygen than ethane and burns with a clean, almost invisible blue flame.
In everyday life: Ethanol burns in spirit burners in the school lab, in flambéed desserts and as bioethanol blended into petrol.
Combustion of octane
2 C₈H₁₈ + 25 O₂ → 16 CO₂ + 18 H₂O
Each octane molecule loses its seven C–C and eighteen C–H bonds, and the O=O bonds of the oxygen molecules break too; the atoms form CO₂ with C=O double bonds and water with O–H bonds. Two octane molecules need as many as 25 oxygen molecules – when there is not enough air, poisonous carbon monoxide or soot form instead of CO₂.
In everyday life: Every time a petrol car drives, octane and similar hydrocarbons in the fuel burn like this in the engine cylinders.
Cellular respiration
C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O
Overall, the C–C, C–H, C–O and O–H bonds of glucose and the O=O bonds of oxygen break, and the atoms end up in CO₂ and water. In the cell this does not happen in one go as in a flame but through many enzyme-controlled steps, most of them in the mitochondria, and part of the released energy is stored in ATP; the app shows only the overall equation.
In everyday life: Every breath you take supplies oxygen for this reaction, and the CO₂ you breathe out is one of its products.
Photosynthesis
6 CO₂ + 6 H₂O → C₆H₁₂O₆ + 6 O₂
Overall, the C=O bonds of carbon dioxide and the O–H bonds of water break, and new C–C, C–H, C–O and O–H bonds are built into glucose, while oxygen atoms from water pair up into O₂. The energy comes from light captured by chlorophyll in the chloroplasts and ends up stored in the bonds of glucose; in reality it is a long chain of enzymatic steps, and the app shows only the overall equation.
In everyday life: Every green leaf makes sugar this way, and the oxygen you breathe was released by plants and algae.
Chlorination of methane
CH₄ + Cl₂ → CH₃Cl + HCl
UV light splits the Cl–Cl bond into two chlorine radicals, which start a chain reaction: one C–H bond in methane breaks, a C–Cl bond forms in its place, and the removed hydrogen forms H–Cl. This is a radical substitution – a hydrogen atom is replaced by a chlorine atom; with more chlorine the substitution continues to CH₂Cl₂, CHCl₃ and CCl₄.
In everyday life: Chloromethane is mainly used to make silicones – for example the sealant around a bathtub or flexible baking moulds.
Hydrogenation of ethene
C₂H₄ + H₂ → C₂H₆
On the surface of the metal catalyst the H–H bond breaks and one of the two bonds of the C=C double bond opens; each carbon atom then forms a new C–H bond, leaving a C–C single bond. This is an addition: two molecules combine into one and nothing is split off, so unsaturated ethene becomes saturated ethane.
In everyday life: The same kind of reaction with a nickel catalyst hardens liquid vegetable oils into solid fats, for example for margarine.
Bromination of ethene
C₂H₄ + Br₂ → C₂H₄Br₂
The Br–Br bond breaks and one of the two bonds of the C=C double bond opens; each carbon atom then binds one bromine atom through a new C–Br bond. Because the orange-brown bromine is used up, the colour disappears – a quick test showing that a substance contains a C=C double bond.
In everyday life: In the school lab, decolourising bromine water is the standard way to tell alkenes from alkanes.
Hydration of ethene
C₂H₄ + H₂O → C₂H₅OH
One of the two bonds of the C=C double bond opens and an O–H bond in the water molecule breaks; one carbon atom gains a hydrogen atom (new C–H bond) and the other an OH group (new C–O bond). This addition turns ethene from crude oil into ethanol and is the main route to synthetic ethanol.
In everyday life: Ethanol made this way is used as a solvent and in disinfectants, cosmetics and paints rather than in drinks.
Esterification of acetic acid
CH₃COOH + C₂H₅OH → CH₃COOC₂H₅ + H₂O
The C–O bond to the OH group of acetic acid and the O–H bond of ethanol break; the OH from the acid and the H from the alcohol leave as a water molecule, and a new C–O bond joins the two parts into the ester ethyl acetate. The reaction is reversible and reaches an equilibrium with hardly any heat change; concentrated sulfuric acid speeds it up and binds water, which shifts the equilibrium towards the ester.
In everyday life: Ethyl acetate smells fruity and is the solvent in some nail polish removers and glues; similar esters give fruits their aroma.
Frequently asked questions
What is Chemical Reactions 3D?
Chemical Reactions 3D is a free interactive simulation from StudyFi Lab. It runs right in your browser, with no sign-up and nothing to install. Watch atoms rearrange as bonds break and form
Is Chemical Reactions 3D free?
Yes. Chemical Reactions 3D is completely free. It works in the browser on computers, tablets and phones, and you do not need an account.
What can I explore in Chemical Reactions 3D?
Formation of water, Synthesis of ammonia, Synthesis of hydrogen chloride, Formation of nitrogen monoxide, Formation of nitrogen dioxide, Oxidation of sulfur dioxide, Combustion of carbon monoxide, Formation of carbonic acid, Decomposition of hydrogen peroxide, Electrolysis of water, Decomposition of ozone, Decomposition of dinitrogen tetroxide, Combustion of methane, Combustion of ethane, Combustion of propane, Combustion of ethanol, Combustion of octane, Cellular respiration, Photosynthesis, Chlorination of methane, Hydrogenation of ethene, Bromination of ethene, Hydration of ethene, Esterification of acetic acid.
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