Question

Although magnesium metal does not react with water at room temperature, it does react vigorously with steam at higher temperatures, releasing elemental hydrogen gas from the water. \(\mathrm{Mg}(s)+2 \mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{Mg}(\mathrm{OH})_{2}(s)+\mathrm{H}_{2}(g)\) Identify which element is being oxidized and which is being reduced.

Short answer

In the chemical reaction \( Mg(s) + 2 H_2O(g) \rightarrow Mg(OH)_2(s) + H_2(g) \), Magnesium (Mg) is being oxidized as its oxidation state increases from 0 to +2, and Hydrogen (H) is being reduced as its oxidation state decreases from +1 to 0.

Step-by-step solution

Write down the given chemical equation and its oxidation states

The given chemical equation is: \( Mg(s) + 2 H_2O(g) \rightarrow Mg(OH)_2(s) + H_2(g) \) First, let's assign the oxidation states to each element. Keep in mind the following oxidation state rules: 1. The oxidation state of an element in its elemental form is 0. 2. The oxidation state of oxygen in a compound is usually -2. 3. The oxidation state of hydrogen in a compound is usually +1. Using these rules, we can assign the oxidation states as follows: \( Mg(0) + 2 H_{(+1)} O_{(-2)} \rightarrow Mg(+2)[ O_{(-2)} H_{(+1)} ]_2 + H_{(+1)} H_{(+1)} \)

Identify the changes in oxidation states

Now, let's identify the changes in oxidation states for each element: 1. Magnesium: Mg(0) -> Mg(+2); it increased by 2. 2. Oxygen: O(-2) -> O(-2); it remained the same. 3. Hydrogen: H(+1) -> H(+1); it remained the same for the oxygen-bound hydrogen atoms and changed for the hydrogen atoms in the gas. To determine which element is oxidized and which is reduced, we need to focus on the elements that have a change in their oxidation states, which are Magnesium and Hydrogen.

Determine which element is oxidized and which is reduced

According to the changes in oxidation states identified in Step 2: 1. Magnesium: Mg(0) -> Mg(+2); the oxidation state increased by 2. This indicates that Magnesium loses 2 electrons and is being oxidized. 2. Hydrogen: H(+1) -> H(0); the oxidation state decreased by 1. This indicates that Hydrogen gains 1 electron and is being reduced. So, in this chemical reaction, the element Magnesium is being oxidized, and the element Hydrogen is being reduced.

Key concepts

Oxidation States

Understanding oxidation states is crucial in redox reactions. The oxidation state, also known as oxidation number, helps identify how many electrons are lost or gained by an atom during a chemical reaction. Here are some basic rules to follow:

• The oxidation state of an element in its natural form is always 0, like magnesium (Mg) in its metallic form.
• Oxygen generally has an oxidation state of -2 in compounds.
• Hydrogen usually has an oxidation state of +1 when bonded with nonmetals.

Identifying the changes in these states tells us which elements are being oxidized or reduced.

Magnesium

Magnesium is a reactive metal that plays a significant role in redox reactions. In our exercise, magnesium starts in its elemental form with an oxidation state of 0. During the reaction with steam, magnesium reacts to form magnesium hydroxide ( Mg(OH)_2) and becomes Mg(+2).
This increase in oxidation state from 0 to +2 indicates that magnesium loses two electrons, undergoing oxidation. Magnesium's ability to donate electrons makes it a frequent participant in redox reactions.

Hydrogen

Hydrogen is another essential player in redox processes. Commonly found as H(+1) in compounds like water ( H_2O), hydrogen can either gain or lose electrons. In the given reaction, hydrogen starts at an oxidation state of +1 when bound to oxygen in water.
When hydrogen gas ( H_2) is formed, its oxidation state reduces to 0, indicating a gain of electrons. The transition from H(+1) to H(0) reveals hydrogen's role as the reduced element in this redox reaction.

Oxidation and Reduction

Oxidation and reduction, often remembered by the mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain), describe electron transfer in reactions. In our example, magnesium undergoes oxidation because it loses two electrons, going from an oxidation state of 0 to +2.
Conversely, hydrogen experiences reduction by gaining electrons, shifting from +1 in water to 0 in hydrogen gas. Redox reactions revolve around such electron exchanges and changes in oxidation states.

Electron Transfer

Electron transfer is at the core of redox reactions. It involves the movement of electrons from one substance to another. Here, magnesium acts as a reducing agent by losing electrons and transferring them to hydrogen. This electron donation from Mg to H facilitates their oxidation and reduction, respectively.
Understanding the flow of electrons helps visualize chemical changes and predict the outcomes of such reactions. Redox reactions often involve this dynamic exchange, driving the transformation of elements.