Ellingham diagrams free#

For example, metallic aluminium can reduce iron oxide to metallic iron, the aluminium itself being oxidized to aluminium oxide.

Ellingham diagrams free#

A reduced substance (such as a metal), whose Gibbs free energy of formation is lower on the diagram at a given temperature, will reduce an oxide whose free energy of formation is higher on the diagram.

According to the Boudouard reaction, carbon monoxide is the dominant oxide of carbon at higher temperatures (above about 700 ☌), and the higher the temperature (above 700 ☌) the more effective a reductant (reducing agent) carbon is.

The formation free energy of carbon dioxide (CO 2) is almost independent of temperature, while that of carbon monoxide (CO) has negative slope and crosses the CO 2 line near 700 ☌.

Highly unstable oxides like AgĢO and HgO easily undergo thermal decomposition.

Stability of metallic oxides decreases with increase in temperature.

For example, the line for Al (oxidation of aluminium) is found to be below that for Fe (formation of Fe

The lower the position of a metal's line in the Ellingham diagram, the greater is the stability of its oxide.

The slope is proportional to ΔS, which is fairly constant with temperature.

Curves in the Ellingham diagrams for the formation of metallic oxides are basically straight lines with a positive slope.

The oxide with the more negative ΔG will be formed and the other oxide will be reduced. If two metals are present, two equilibria have to be considered. The blue line for the formation of CO 2 is approximately horizontal, since the reaction C(s) + OĢ(g) → CO 2(g) leaves the number of moles of gas unchanged so that ΔS is small.Īs with any chemical reaction prediction based on purely thermodynamic grounds, a spontaneous reaction may be very slow if one or more stages in the reaction pathway have very high activation energies E A. At a sufficiently high temperature, the sign of ΔG may invert (becoming positive) and the oxide can spontaneously reduce to the metal, as shown for Ag and Cu.įor oxidation of carbon, the red line is for the formation of CO: C(s) + 1⁄ 2 OĢ(g) → CO(g) with an increase in the number of moles of gas, leading to a positive ΔS and a negative slope. Since these reactions are exothermic, they always become feasible at lower temperatures.

The slope of the plots dΔG/dT = − ΔS is therefore positive for all metals, with ΔG always becoming more negative with lower temperature, and the lines for all the metal oxides are approximately parallel. For the oxidation of each metal, the dominant contribution to the entropy change (ΔS) is the removal of 1⁄ 2 mol OĢ, so that ΔS is negative and roughly equal for all metals. In the temperature ranges commonly used, the metal and the oxide are in a condensed state (liquid or solid), and oxygen is a gas with a much larger molar entropy. The diagram at right refers to 1 mole O, so that for example the line marked Crģ shows ΔG for the reaction 2/3 Cr(s) + 1⁄ 2 Oģ(s), which is 1⁄ 3 of the molar Gibbs energy of formation ΔG f°( Cr For comparison of different reactions, all values of ΔG refer to the reaction of the same quantity of oxygen, chosen as one mole O ( 1⁄ 2 mol OĢ by others. The Ellingham diagram plots the Gibbs free energy change (ΔG) for each oxidation reaction as a function of temperature. Simple Ellingham diagram for high temperature (0☌ – 2500☌) oxidation of several metals and carbon