burning hydrocarbons experiment

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Combustion Reaction Definition and Examples

A combustion reaction is an exothermic chemical reaction between a fuel and an oxidizer that forms an oxidized product. In general chemistry, it is one of the main types of chemical reactions . Combustion is a reaction between a hydrocarbon fuel (e.g., coal, propane, wood, methane) and molecular oxygen (O 2 ), producing carbon dioxide (CO 2 ), water (H 2 O), and heat. Heat provides the activation energy to start the chemical reaction. The combination of oxygen, fuel, and heat forms the fire triangle , which is one way to represent the requirements for combustion.

General Form of Combustion Reaction Equation

The general form of a combustion reaction is:

hydrocarbon + oxygen → carbon dioxide + water + heat

C x H y + O 2 → CO 2 + H 2 O

Examples of Combustion Reactions

Combustion is also called burning. So, any example of burning you can think of is a combustion reaction, including burning matches, candles, campfires, and gas burners. Here are examples of balanced equations for combustion reactions:

• Combustion of methane CH 4 (g) + 2 O 2 (g) → CO 2 (g) + 2 H 2 O(g)
• Burning of naphthalene C 10 H 8  + 12 O 2  → 10 CO 2  + 4 H 2 O
• Combustion of ethane 2 C 2 H 6  + 7 O 2  → 4 CO 2  + 6 H 2 O
• Combustion of butane (commonly found in lighters) 2C 4 H 10 (g) +13O 2 (g) → 8CO 2 (g) +10H 2 O(g)
• Combustion of methanol (also known as wood alcohol) 2CH 3 OH(g) + 3O 2 (g) → 2CO 2 (g) + 4H 2 O(g)
• Combustion of propane (used in gas grills, fireplaces, and some cookstoves) 2C 3 H 8 (g) + 7O 2 (g) → 6CO 2 (g) + 8H 2 O(g)

How to Recognize a Combustion Reaction

You’ll know you have a combustion reaction when you see a hydrocarbon (molecule containing carbon and hydrogen) and oxygen gas (O 2 ) on the reactant side (left side) of the reaction arrow and carbon dioxide (CO 2 ) and water (H 2 O) on the product side (right side) of the reaction arrow. Also, combustion using oxygen always produces heat. The reaction still requires activation energy to start, but more heat is released by combustion than is absorbed starting it.

Many combustion reaction produce flames. If you see fire, it indicates a combustion reaction. However, combustion often occurs without fire. For example, smoldering is combustion without flames.

Sometimes it’s harder to recognize a combustion reaction because the reactant contains its own oxidizer (oxygen) or because combustion is incomplete, forming other products besides carbon dioxide and water. For example, some rockets rely on the reaction between Aerozine 50 (C 2 H 12 N 4 ) and nitrogen tetroxide (N 2 O 4 ). If you’re clever, you’ll see Aerozine 50 contains the necessary chemical bonds to act as a fuel (carbon-hydrogen and carbon-nitrogen) and nitrogen tetroxide supplies oxygen for combustion.

Then, there are forms of combustion that don’t even involve oxygen.

Combustion Without Oxygen

Technically, oxidation doesn’t always require oxygen , so combustion can occur without oxygen, too.

An oxidizer accepts electrons, usually by supplying oxygen to a chemical reaction. Other oxidizers include the halogens (fluorine, chlorine, etc.). Metallic fuels burn using fluoropolymers (e.g., Teflon, Viton), without the need for any oxygen at all.

Complete Versus Incomplete Combustion

Like other chemical reactions, combustion is subject to a limiting reactant and doesn’t always proceed to completion.

• Complete combustion or “clean combustion” occurs when the oxidation of a hydrocarbon only produces carbon dioxide and water. Burning candle wax is a good example of complete combustion. Heat from the burning wick vaporizes wax (a hydrocarbon). Wax reacts with oxygen, releasing carbon dioxide and water. The wax burns away and the carbon dioxide and water dissipate into the air.
• Incomplete combustion or “dirty combustion” is incomplete hydrocarbon oxidation, producing carbon monoxide (CO), carbon (soot), and other products, in addition to carbon dioxide and water. Wood and most fossil fuels undergo incomplete combustion, releasing these additional waste products.
• Lackner, Maximilian; Winter, Franz; Agarwal, Avinash K., eds. (2010). Handbook of Combustion . Wiley-VCH. ISBN 978-3-527-32449-1.
• Law, C.K. (2006). Combustion Physics . Cambridge University Press. ISBN 9780521154215.
• Schmidt-Rohr, K (2015). “Why Combustions Are Always Exothermic, Yielding About 418 kJ per Mole of O 2 “.  J. Chem. Educ .  92  (12): 2094–2099.  doi: 10.1021/acs.jchemed.5b00333

Related Posts

Burning hydrocarbons

I can write balanced symbol equations for combustion reactions, and describe the difference between complete and incomplete combustion.

Lesson details

Key learning points.

• Combustion is the chemical combination of a substance with oxygen, involving the production of heat and light.
• Complete combustion of fuels (like hydrocarbons) involve the production of carbon dioxide and water.
• The chemical test for oxygen is a combustion reaction, involving relighting a glowing splint.
• Incomplete combustion produces carbon particulates (soot) and carbon monoxide due to reduced oxygen levels.

Common misconception

All combustion is complete, producing only CO₂ and water.

Highlight the differences between complete and incomplete combustion.

Combustion - Combustion is an exothermic reaction where a substance reacts with oxygen.

Complete combustion - The complete combustion of fuels, such as hydrocarbons, occurs when enough oxygen is present to convert all the fuel into carbon dioxide and water.

Hydrocarbons - Hydrocarbons are compounds containing only hydrogen and carbon atoms.

Incomplete combustion - Incomplete combustion is when a substance reacts only partially with oxygen, such as when carbon burns in air, producing carbon dioxide, carbon monoxide and soot (unburnt carbon).

This content is © Oak National Academy Limited ( 2024 ), licensed on Open Government Licence version 3.0 except where otherwise stated. See Oak's terms & conditions (Collection 2).

Starter quiz

6 questions.

hydrocarbons -  

molecules made of hydrogen and carbon atoms only

conservation of mass -  

the total mass of reactants equals the total mass of products formed

exothermic reactions -  

chemical reactions in which thermal energy is given out

fossil fuels -  

fuels made over millions of years from dead plants and animals

Burning Hydrocarbons

The equations for the combustion (burning) of hydrocarbons depend on the amount of oxygen that is available.

Burning with plenty of oxygen

• The release of energy.
• The oxidation of carbon and hydrogen to create carbon dioxide and water.

Burning with a lack of oxygen

• 2C + O 2 → 2CO

Heavy Fractions of Crude Oil

Heavy fractions of crude oil do NOT make good fuels because they:

Don't ignite easily

Have high boiling points

Have low volatility

1 Atomic Structure

1.1 Atoms & Elements

1.1.1 Elements & Compounds

1.1.2 Chemical Reactions & Equations

1.1.3 Mixtures

1.1.4 Separating Mixtures

1.1.5 Model of the Atom

1.1.6 Atom Size & Number

1.1.7 Atomic Number - Calculations

1.1.8 Isotopes

1.1.9 Isotopes HyperLearning

1.1.10 Isotopes - Calculations

1.1.11 Relative Atomic Mass

1.1.12 Electron Structure

1.1.13 Periodic Table

1.1.14 Metals & Non-Metals

1.1.15 Noble Gases

1.1.16 Halogens

1.1.17 Alkali Metals

1.1.18 Transition Metals

1.1.19 End of Topic Test - Atomic Structure

1.1.20 Grade 9 - Atoms & Elements

1.1.21 Exam-Style Questions - Atomic Structure

1.1.22 Application Questions - Atomic Structure

2 Chemical Bonding

2.1 Chemical Bonds

2.1.1 Types of Bonds

2.1.2 Ionic Bonds

2.1.3 Ionic Compounds

2.1.4 Covalent & Metallic Bonds

2.1.5 Representing Covalent Bonds

2.1.6 Dots-And-Cross Diagrams

2.1.7 Exam-Style Questions - Chemical Bonds

2.2 States of Matter

2.2.1 States of Matter

2.2.2 Changing State

2.2.3 Changing State HyperLearning

2.2.4 State Symbols

2.2.5 Exam-Style Questions - States of Matter

2.3 Chemical Properties

2.3.1 Chemical Properties of Ionic Compounds

2.3.2 Chemical Properties of Small Molecules

2.3.3 Molecular Forces & Polymers

2.3.4 Metals & Giant Covalent Structures

2.3.5 Alloys & Conductors

2.3.6 Carbon Structures - Diamond & Graphite

2.3.7 Carbon Structures - Graphene & Fullerenes

2.3.8 Carbon Structures - Summary

2.3.9 Nanoparticles

2.3.10 Nanotechnology

2.3.11 Surface Area to Volume Ratio & Nanotechnology Uses

2.3.12 End of Topic Test - Chemical Bonding

2.3.13 Grade 9 - Bonding & Structure

2.3.14 Exam-Style Questions - Allotropes of Carbon

3 Quantitative Chemistry

3.1 Chemical Measurements

3.1.1 Conservation of Mass

3.1.2 Relative Formula Mass

3.1.3 Relative Formula Mass - Calculations

3.1.4 Measuring Mass

3.1.5 Equations & Formula

3.1.6 Moles 1

3.1.7 Moles 2

3.1.8 Moles 3

3.1.9 Moles - Calculations

3.1.10 Percentage Yield

3.1.11 Percentage Yield - Calculations

3.1.12 Moles & Yield

3.1.13 Atom Economy

3.1.14 Atom Economy - Calculations

3.1.15 Concentrations & Amounts

3.1.16 Amount of Substance in Relation to Volumes and Gas

3.1.17 Concentrations & Amounts - Calculations

3.1.18 Empirical Formula - Calculations

3.1.19 End of Topic Test - Quant Chemistry

3.1.20 Grade 9 - Quantitative Chemistry

3.1.21 Exam-Style Questions - Moles

4 Chemical Changes

4.1 Acids, Alkalis & Redox Reactions

4.1.1 Acids & Alkalis

4.1.2 Redox Reactions

4.2 Reactivity of Metals

4.2.1 Reactivity Diagram

4.2.2 Displacement Reactions

4.2.3 Reactions with Water & Acid

4.2.4 Reactivity Tests

4.2.5 Extraction of Metals

4.2.6 Exam-Style Questions - Displacement Reactions

4.3 Reactions of Acids

4.3.1 Reactions of Metals with Acids

4.3.2 Neutralisation of Acids

4.3.3 Soluble Salts

4.3.4 Required Practical: Separating Mixtures

4.3.5 The pH Scale

4.3.6 Measuring pH

4.3.7 Neutralisation & Titrations 1

4.3.8 Neutralisation & Titrations 2

4.3.9 Strong and Weak Acids

4.4 Electrolysis

4.4.1 Electrolysis

4.4.2 Metal Extraction

4.4.3 Electrolysis of Aqueous Solutions & Half Equations

4.4.4 End of Topic Test - Chemical Changes

4.4.5 Grade 9 - Chemical Changes

4.4.6 Exam-Style Questions - Electrolysis

5 Energy Changes

5.1 Exothermic & Endothermic Reactions

5.1.1 Energy Conservation

5.1.2 Exothermic & Endothermic Reactions

5.1.3 Reaction Profiles

5.1.4 Reaction Profiles 2

5.2 Chemical Cells & Fuel Cells

5.2.1 Chemical Cells

5.2.2 Batteries & Charging of Cells

5.2.3 Fuel Cells

5.2.4 Hydrogen Fuel Cell

5.2.5 End of Topic Test - Energy Changes

5.2.6 Grade 9 - Energy Changes

5.2.7 Exam-Style Questions - Fuel Cells

6 The Rate & Extent of Chemical Change

6.1 Rate of Reaction

6.1.1 Chemical Reactions

6.1.2 Collision Theory

6.1.3 Catalysts

6.1.4 Catalysts in Industry

6.2 Reversible Reactions

6.2.1 Reversible Reactions

6.2.2 Equilibrium

6.2.3 Equilibrium Position

6.2.4 Equilibrium Position 2

6.2.5 End of Topic Test - Rates of Reaction

6.2.6 Grade 9 - Rate & Extent of Chemical Change

6.2.7 Exam-Style Questions - Reversible Reactions

7 Organic Chemistry

7.1 Carbon Compounds

7.1.1 Crude Oil

7.1.2 Alkanes

7.1.3 Fractional Distillation

7.1.4 Homologous Series

7.1.5 Burning Hydrocarbons

7.1.6 Cracking

7.1.7 Uses of Alkenes

7.1.8 Exam-Style Questions - Organic Compounds

7.2 Alkenes & Alcohols

7.2.1 Structure of Alkenes

7.2.2 Reactions of Alkenes

7.2.3 End of Topic Test - Alkanes & Alkenes

7.2.4 Alcohols

7.2.5 Reactions of Alcohols

7.2.6 Carboxylic Acids

7.2.7 Reactions of Carboxylic Acids

7.2.8 Addition Polymers

7.2.9 Condensation Polymers

7.2.10 Naturally Occurring Polymers - Proteins

7.2.11 Naturally Occurring Polymers - DNA & Sugars

7.2.12 End of Topic Test - Alcohols, Acids & Polymers

7.2.13 Grade 9 - Organic Chemistry

7.2.14 Exam-Style Questions - Alcohols

8 Chemical Analysis

8.1 Purity, Formulations & Chromatography

8.1.1 Purity & Formulations

8.1.2 Chromatography

8.1.3 Paper Chromatography 1

8.1.4 Paper Chromatography 2

8.1.5 Required Practical - Paper Chromatography

8.2 Identification of Common Gases

8.2.1 Identification of Hydrogen & Oxygen

8.2.2 Identification of Carbon Dioxide & Chlorine

8.3 Identifying Ions

8.3.1 Flame Tests

8.3.2 Required Practical: Flame Tests

8.3.3 Metal Hydroxides

8.3.4 Carbonates

8.3.5 Halides, Sulfates

8.3.6 Methodology for Chemical Analysis

8.3.7 Flame Emission Spectroscopy

8.3.8 End of Topic Test - Chemical Analysis

8.3.9 Grade 9 - Chemical Analysis

8.3.10 Exam-Style Questions - Identification of ions

9 Chemistry of the Atmosphere

9.1 The Earth's Atmosphere

9.1.1 Proportions of Gases in the Atmosphere

9.1.2 Oxygen & Carbon Dioxide in the Atmosphere

9.1.3 Greenhouse Gases & Human Activities

9.1.4 Greenhouse Gases & Human Activities

9.1.5 Greenhouse Gases & Climate Change

9.1.6 Carbon Footprints

9.1.7 Common Atmospheric Pollutants & Their Properties

9.1.8 End of Topic Test - The Earth

9.1.9 Grade 9 - Chemistry of the Atmosphere

9.1.10 Exam-Style Questions - Earth's Atmosphere

10 Using Resources

10.1 Using the Earth's Resources

10.1.1 Potable Water & its Sources

10.1.2 Potable Water in the UK & in Saudi Arabia

10.1.3 Natural Resources

10.1.4 Waste Water

10.1.5 Metal Ore Extraction & the Properties of Copper

10.1.6 Overexploitation & Alternative Extraction Methods

10.2 Life Cycle Assessments & Recycling

10.2.1 Life Cycle Assessments

10.2.2 Difficulties with LCA

10.2.3 Reducing Resource Usage

10.2.4 Reducing Use of Metal and Glass

10.2.5 Exam-Style Questions - Life Cycle Assessments

10.3 Using Materials

10.3.1 Corrosion

10.3.2 Alloys

10.3.3 Ceramics

10.3.4 Polymers & Composites

10.4 The Haber Process & NPK Fertilisers

10.4.1 The Haber Process

10.4.2 The Haber Process - Equation & Trade-Offs

10.4.3 Formulated vs Manure Fertilisers

10.4.4 NPK Fertilisers 1

10.4.5 NPK Fertilisers 2

10.4.6 End of Topic Test - Resources

10.4.7 Grade 9 - Using Resources

10.4.8 Exam-Style Questions - The Haber Process

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Combustion of a Hydrocarbon AKA: Burning Candle

A candle is burned in a gas jar and its behaviour observed.

• Gas jar and lid
• Candle in tray
• Dry blue cobalt chloride paper
• Heatproof mat

Related Links

• Burning candle apparatus

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Comparing heat energy from burning alcohols

In association with Nuffield Foundation

• Four out of five

In this investigation, students use a spirit burner to burn various alcohols while measuring and comparing the amount of heat energy produced

This experiment is suitable for pre-16 students, possibly as an introduction to a topic on fuels. It can be taken further if used with post-16 students who can calculate values for enthalpy changes of combustion, with subsequent discussion about heat losses and incomplete combustion.

The alcohols should be provided in labelled spirit burners ready to use. If each group investigates one alcohol, the experiment can be done in around 20 minutes. It is better if each spirit burner is used by more than one group of students. Variation of results will add substance to a discussion about errors.

• Eye protection
• Retort stand and clamp
• Conical flask, 150 cm 3 or larger
• Measuring cylinder, 100 cm 3
• Thermometer (–10 °C to +110 °C)
• Access to balances, preferably several, to avoid queuing
• Access to spirit burners with wicks and caps, containing the alcohols listed (note 1)

Apparatus notes

• Suitable spirit burners are hard to come by. Ideally they should be small, with a capacity of 50 cm 3 or less. Pictures and information in suppliers’ catalogues can be misleading. It is important that the wick fits tightly in the wick holder and that the wick holder fits tightly in the burner. If capacity is more than 50 cm 3 , reduce it, for instance by packing with mineral wool, or partially filling with epoxy. Refer to CLEAPSS L195 ’Safer chemicals, safer reactions’. One possible source is: A.J.Cope & Son Ltd , Unit 10, Cliffside Trade Park, Motherwell Way, Grays, Essex, RM20 3XD.
• Methanol 
• Ethanol 
• Propan-1-ol 
• Propan-2-ol 
• Butan-1-ol 

Health, safety and technical notes

• Read our standard health and safety guidance .
• Wear eye protection throughout.
• Methanol, CH 3 OH(l), (HIGHLY FLAMMABLE, TOXIC) – see CLEAPSS Hazcard HC040b . Methanol is volatile and has a low flash point.
• Ethanol, CH 3 CH 2 OH(l), (HIGHLY FLAMMABLE) – see CLEAPSS Hazcard HC040A . Ethanol is volatile and has a low flash point.
• Propan-1-ol, CH 3 CH 2 CH 2 OH(l), (HIGHLY FLAMMABLE, IRRITANT, HARMFUL) – see CLEAPSS Hazcard HC084A . Propan-1-ol is volatile and has a low flash point.
• Propan-2-ol, CH 3 CHOHCH 3 (l), (HIGHLY FLAMMABLE, IRRITANT, HARMFUL) – see CLEAPSS Hazcard HC084A. Propan-2-ol is volatile and has a low flash point.
• Butan-1-ol, CH3CH 2 CH 2 CH 2 OH(l), (HIGHLY FLAMMABLE, IRRITANT, HARMFUL) – see CLEAPSS Hazcard HC084B . Butan-1-ol is volatile and has a low flash point.

Source: Royal Society of Chemistry

Equipment required for measuring heat energy from burning alcohol.

• Measure 100 cm 3  of cold tap water into a conical flask.
• Clamp the flask at a suitable height so that a spirit burner can easily be placed below.
• Weigh the spirit burner (and cap) containing the alcohol and record this mass and the name of the alcohol.
• Record the initial temperature of the water in the flask.
• Place the spirit burner under the flask and light the wick.
• Allow the alcohol to heat the water so the temperature rises by about 40 °C.
• Replace the cap to extinguish the flame.
• Reweigh the spirit burner and cap, and record this mass.
• Work out the mass of alcohol used.
• Using a fresh 100 cm 3  of cold tap water, repeat the experiment with another alcohol.

Teaching notes

Get the class to record and share the results. Do not be surprised if groups get different answers for a given alcohol. Heat losses will almost certainly vary considerably.

Subsequent discussion depends on the level of the students’ experience.

Student questions

Here are some possible questions to ask students:

• Which alcohol produces the most energy per gram?
• Which alcohol produces the most energy per mole?
• Write equations for the complete combustion of each alcohol.
• Propan-1-ol and propan-2-ol are isomers (same molecular formula, different structures). Do they produce the same amount of heat on combustion?
• Does all the heat produced by combustion go into raising the temperature of the water?
• Is it possible that combustion may be incomplete, giving carbon monoxide among the products?
• Alcohols can be used as a substitute for hydrocarbon fuels, and so methods of producing alcohols are very important. What process converts sugar into alcohol and carbon dioxide?

Notes on questions

• On question 6, stress the dangers accompanying the production of carbon monoxide.

More resources

Add context and inspire your learners with our short career videos showing how chemistry is making a difference .

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

© Nuffield Foundation and the Royal Society of Chemistry

• 14-16 years
• 16-18 years
• Practical experiments
• Thermodynamics
• Quantitative chemistry and stoichiometry

Specification

• 9.28C Investigate the temperature rise produced in a known mass of water by the combustion of the alcohols ethanol, propanol, butanol and pentanol
• The heat change, q, in a reaction is given by the equation q = mcΔT; where m is the mass of the substance that has a temperature change ΔT and a specific heat capacity c.
• Students should be able to: use this equation to calculate the molar enthalpy change for a reaction.
• e) determination of enthalpy changes directly from appropriate experimental results, including use of the relationship: q = mcΔT
• 5 i. understand experiments to measure enthalpy changes in terms of: processing results using the expression: energy transferred = mass x specific heat capacity × temperature change (Q=mcΔT)
• In combustion, a substance reacts with oxygen releasing energy.
• Fuels burn releasing different quantities of energy.
• The quantity of heat energy released can be determined experimentally and calculated using, Eₕ = cmΔT.
• PRACTICAL: Determination of an enthalpy change of combustion
• (f) the combustion reactions of hydrocarbons and other fuels
• (g) how to determine experimentally the energy per gram released by a burning fuel
• (v) the uses of ethanol as a solvent and as a fuel and the social, economic and environmental factors that affect the development of bioethanol fuel
• determine the enthalpy changes for combustion and neutralisation using simple apparatus; and
• 2.8.6 recall experimental methods to determine enthalpy changes;
• 2.8.7 calculate enthalpy changes from experimental data using the equation q = mcΔT;

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Hydrogen co-combustion of aromatic volatile organic compounds over pd/al 2 o 3 catalyst.

Graphical Abstract

1. Introduction

2. results and discussion, 2.1. structure evolution of pd/al 2 o 3 catalyst, 2.2. catalytic performance test, 3. experimental section, 3.1. materials, 3.2. catalyst preparation, 3.3. characterization, 3.4. reactor setup and performance test, 4. conclusions, supplementary materials, author contributions, data availability statement, conflicts of interest.

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Ullah, L.; Munsif, S.; Cao, L.; Murthy, P.R.; Zhang, J.-C.; Li, W.-Z. Hydrogen Co-Combustion of Aromatic Volatile Organic Compounds over Pd/Al 2 O 3 Catalyst. Catalysts 2024 , 14 , 563. https://doi.org/10.3390/catal14090563

Ullah L, Munsif S, Cao L, Murthy PR, Zhang J-C, Li W-Z. Hydrogen Co-Combustion of Aromatic Volatile Organic Compounds over Pd/Al 2 O 3 Catalyst. Catalysts . 2024; 14(9):563. https://doi.org/10.3390/catal14090563

Ullah, Lutf, Sehrish Munsif, Long Cao, Palle Ramana Murthy, Jing-Cai Zhang, and Wei-Zhen Li. 2024. "Hydrogen Co-Combustion of Aromatic Volatile Organic Compounds over Pd/Al 2 O 3 Catalyst" Catalysts 14, no. 9: 563. https://doi.org/10.3390/catal14090563

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