Chemistry · Chemical reactions · Grade 7

Combustion, flames and oxides

Why does a candle go out under a glass, and which part of a flame is hottest?

The situation

A winter evening: in the kitchen a gas hob burns with a blue flame, and on the table there is a candle. Someone covers the candle with a glass tumbler — after a few seconds the flame shrinks and goes out: has the oxygen run out? Which part of a candle flame is hottest, and why is the flame of a gas hob blue while a candle's is yellow? What oxides form when charcoal, sulfur, phosphorus, iron and magnesium burn in oxygen, and what kind of medium do they give in water?

Topic

Combustion, flames and oxides. Gas hob and candle flame: the fire triangle, flame zones and firefighting

Key ideas

Fire triangle
Flame zones
Oxides

Formulas

combustion = fuel + oxygen + ignition temperature

The fire triangle: if any one of the three is removed, the fire goes out.

CH_{4} + 2 O_{2} \overset{}{\to} CO_{2} + 2 H_{2} O

Combustion of natural gas (methane): $Δ H$ = −890 kJ/mol, that is ≈ 55.5 MJ/kg of heat.

C_{25} H_{52} + 38 O_{2} \overset{}{\to} 25 CO_{2} + 26 H_{2} O

Combustion of a paraffin candle (approximate formula C₂₅H₅₂).

t = \frac{V (x_{0} - x_{\min})}{r}

How long a candle burns under a tumbler: $x_{0}$ = 20.95 %, $x_{\min}$ ≈ 16 %, $r$ — the candle's O₂ consumption (mL/s).

P_{2} O_{5} + 3 H_{2} O \overset{}{\to} 2 H_{3} PO_{4}

Acidic oxide + water → acid; basic oxide + water → base: $MgO + H_{2} O \overset{}{\to} Mg (OH)_{2}$ .

Theory

Combustion and the fire triangle

Combustion is the rapid combination of a substance with oxygen, releasing heat and light. Three conditions are needed for combustion: fuel, oxygen (air) and heating the fuel to its ignition temperature. Remove any one and the fire goes out. A candle under a tumbler does not use up all the oxygen: when the O₂ in the air falls to ≈ 16 %, the flame can no longer sustain itself and goes out (this is the limit in the model of this scene; in experiments 15–17 %). A candle burns ≈ 0.1 g of paraffin per minute — that is ≈ 4.5 mL/s of oxygen. A cold copper coil draws heat from the flame — the paraffin vapour cools below its ignition temperature and the fire goes out. If the wick is pinched, the fuel (paraffin vapour) stops arriving.

Practical: structure of a flame

A candle flame has three zones. The inner dark zone is paraffin vapour around the wick, which does not burn yet (≈ 600 °C: evaporation absorbs heat). The middle bright zone lacks oxygen and combustion is incomplete: glowing tiny particles of soot (carbon) give off yellow light (≈ 1000–1200 °C); if a cold porcelain dish is held in it, soot is deposited. The outer zone is thin and almost invisible; there is enough oxygen from the air, combustion is complete and this is the hottest place (≈ 1400 °C). A stick passed horizontally through the flame is scorched at its two edges, while the middle stays almost whole. In a burner, if the air hole is closed the flame is yellow and sooty; if it is opened, the gas mixes with air beforehand: a blue flame with a cooler inner cone (unburnt mixture) inside. The hottest point is slightly above the tip of the inner cone (1200–1600 °C in sources; ≈ 1500 °C in the model). Vessels are heated in this upper part of the flame.

Gas hob: burning natural gas

The natural gas in our homes is mostly methane: $CH_{4} + 2 O_{2} \overset{}{\to} CO_{2} + 2 H_{2} O$ ; when 1 mol of methane burns, 890 kJ (≈ 55.5 MJ per kg) of heat is released. The blue flame of the hob shows that the gas is mixing well with air. If the flame is yellowish-red and smoky, there is not enough air: soot and carbon monoxide CO form — a colourless, odourless and very poisonous gas. That is why a room where gas is burning is ventilated, and if you smell gas, you do not strike a match: you close the tap and open the windows.

Practical: burning substances in oxygen

The burning substance is lowered on a spoon into a jar of oxygen. Charcoal glows white-yellow: $C + O_{2} \overset{}{\to} CO_{2}$ . Sulfur burns with a bright blue-violet flame: $S + O_{2} \overset{}{\to} SO_{2}$ (a gas with a pungent smell). Red phosphorus gives a dazzling flame and thick white smoke: $4 P + 5 O_{2} \overset{}{\to} 2 P_{2} O_{5}$ . Iron wire burns throwing off sparks: $3 Fe + 2 O_{2} \overset{}{\to} Fe_{3} O_{4}$ . Magnesium burns with a brilliant white light: $2 Mg + O_{2} \overset{}{\to} 2 MgO$ . A little water is poured into the jar beforehand — so that hot drops of iron scale do not crack the glass. Then the water is swirled and an indicator is added.

Oxides: names and classes

An oxide is a substance made of two elements, one of which is oxygen (oxidation state −2). The name: element name + (valency) + “oxide”: CO₂ — carbon(IV) oxide, SO₂ — sulfur(IV) oxide, P₂O₅ — phosphorus(V) oxide, MgO — magnesium oxide. Basic oxides are the oxides of most metals (Na₂O, CaO, MgO, CuO): they give salts with acids, and those that dissolve in water form alkalis. Acidic oxides are mostly oxides of non-metals (CO₂, SO₂, SO₃, P₂O₅, SiO₂): they give salts with bases, and most form acids with water (SiO₂ does not react with water). Amphoteric oxides (ZnO, Al₂O₃) react with both acids and alkalis. Neutral (non-salt-forming) oxides: CO, NO, N₂O. In the experiment, a solution of CO₂, SO₂ or P₂O₅ turns litmus red, while MgO turns phenolphthalein raspberry-pink (Mg(OH)₂, pH ≈ 10); Fe₃O₄ does not dissolve in water.

History, firefighting, safety and the model

In the 1770s Antoine Lavoisier burned metals and phosphorus in a closed vessel and weighed them: the mass of the substance increases because part of the air — oxygen — combines with it. This created the oxygen theory of combustion and refuted the old idea of “phlogiston”. In firefighting one side of the triangle is removed: water cools, sand, felt and a carbon dioxide extinguisher block the oxygen, and shutting the gas tap stops the fuel arriving. Water is not sprayed on burning fat or on a live electrical appliance. Safety: safety goggles; tie back long hair; sulfur and phosphorus are burned in a fume cupboard, and only red phosphorus is used; do not stare at the light of magnesium. Model: the candle consumes O₂ at a constant rate and goes out when O₂ is 16 %; the gas in the tumbler is fully mixed; flame temperatures are an approximate smooth model; the pH of the oxide solutions is calculated with Henry's law and constants at 25 °C, with the solutions treated as ideal.

Did you know?

In the 1770s Antoine Lavoisier heated metals in a closed vessel and showed that their mass increases because of the oxygen they take from the air, and he created the oxygen theory of combustion. He also gave oxygen its name, “Oxygène”.
Michael Faraday gave Christmas lectures on the candle for children at the Royal Institution in London in 1848 and 1860; the 1860 lectures were published in 1861 as the book “The Chemical History of a Candle”, and it is still one of the most famous popular chemistry books.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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Reversible reactions and equilibrium — Grade 8
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