Hello Guys, In this Video we are going to discuss a very important topics in metallurgy for Jee mains, Advanced, NEET, AIIMS and Olympiads. Its Ellingham Diagram , also called ED in Pyrometallurgy.
The Gibbs free energy (∆G) of a reaction is a measure of the thermodynamic driving force that makes a reaction occur. A negative value for ∆G indicates that a reaction can proceed
spontaneously without external inputs, while a positive value indicates that it will not. The equation for Gibbs free energy is:
∆G = ∆H - T∆S
where ∆H is the enthalpy, T is absolute temperature, and ∆S is entropy.
The enthalpy (∆H) is a measure of the actual energy that is liberated when the reaction occurs (the“heat of reaction”). If it is negative, then the reaction gives off energy, while if it is positive the
reaction requires energy.
The entropy (∆S) is a measure of the change in the possibilities for disorder in the products compared to the reactants. For example, if a solid (an ordered state) reacts with a liquid (a somewhat less ordered state) to form a gas (a highly disordered state), there is normally a large positive change in the entropy for the reaction
Construction of an Ellingham Diagram
An Ellingham diagram is a plot of ∆G versus temperature. Since ∆H and ∆S are essentially constant with temperature unless a phase change occurs, the free energy versus temperature plot can be drawn as a series of straight lines, where ∆S is the slope and ∆H is the y-intercept. The slope of the line changes when any of the materials involved melt or vaporize. Free energy of formation is negative for most metal oxides, and so the diagram is drawn with
∆G=0 at the top of the diagram, and the values of ∆G shown are all negative numbers. Temperatures where either the metal or oxide melt or vaporize are marked on the diagram.
The Ellingham diagram shown is for metals reacting to form oxides (similar diagrams can also be drawn for metals reacting with sulfur, chlorine, etc., but the oxide form of the diagram is most common). The oxygen partial pressure is taken as 1 atmosphere, and all of the reactions are normalized to consume one mole of O2.
The majority of the lines slope upwards, because both the metal and the oxide are present as condensed phases (solid or liquid). The reactions are therefore reacting a gas with a condensed
phase to make another condensed phase, which reduces the entropy. A notable exception to this is the oxidation of solid carbon. The line for the reaction .
C+O2 ==CO2
is a solid reacting with a mole of gas to produce a mole of gas, and so there is little change in
entropy and the line is nearly horizontal. For the reaction
2C+O2 == 2CO
we have a solid reacting with a gas to produce two moles of gas, and so there is a substantial
increase in entropy and the line slopes rather sharply downward. Similar behavior can be seen in parts of the lines for lead and lithium, both of which have oxides that boil at slightly lower temperatures than the metal does.
There are three main uses of the Ellingham diagram:
1. Determine the relative ease of reducing a given metallic oxide to metal;
2. Determine the partial pressure of oxygen that is in equilibrium with a metal oxide at a given temperature; and
3. Determine the ratio of carbon monoxide to carbon dioxide that will be able to reduce the oxide to metal at a given temperature.
Ease of Reduction The position of the line for a given reaction on the Ellingham diagram shows the stability of the oxide as a function of temperature. Reactions closer to the top of the diagram are the most “noble” metals (for example, gold and platinum), and their oxides are unstable and easily reduced. As we move down toward the bottom of the diagram, the metals become progressively more reactive and their oxides become harder to reduce. A given metal can reduce the oxides of all other metals whose lines lie above theirs on the diagram. For example, the 2Mg + O2 == 2MgO line lies below the Ti + O2 == TiO2 line, and so magnesium can reduce titanium oxide to metallic titanium. Since the 2C + O2 == 2CO line is downward-sloping, it cuts across the lines for many of the other metals. This makes carbon unusually useful as a reducing agent, because as soon as the carbon oxidation line goes below a metal oxidation line, the carbon can then reduce the metal oxide to metal. So, for example, solid carbon can reduce chromium oxide once the temperature exceeds approximately 1225°C, and can even reduce highly-stable compounds like silicon dioxide and titanium dioxide at temperatures above about 1620°C and 1650°C, respectively. For less stable oxides, carbon monoxide is often an adequate reducing agent.
Other many important topics about ED is discussed in the lecture.
Detailed Ellingham Diagram Link:-
drive.google.c...
Image credit:- Tata Steels Limited
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