Question

For the reaction $\mathrm{Mg}(s)+{\mathrm{Cl}}_2(g)\to {\mathrm{MgCl}}_2(s)$, illustrate how electrons are gained and lost during the reaction.

Short answer

In the redox reaction $\mathrm{Mg}(s)+{\mathrm{Cl}}_2(g)\to {\mathrm{MgCl}}_2(s)$, magnesium loses 2 electrons, changing its oxidation state from 0 to +2 (${\mathrm{Mg}}^0\to {\mathrm{Mg}}^{2+}+2e^{-}$). Each chlorine atom gains 1 electron, changing its oxidation state from 0 to -1 ($2\times ({\mathrm{Cl}}^0+e^{-}\to {\mathrm{Cl}}^{-})$). The electrons are transferred from magnesium to chlorine during the reaction, and the resulting ions form magnesium chloride, ${\mathrm{MgCl}}_2$.

Step-by-step solution

Identify the oxidation states of each element in the reactants and products

For the reaction: $\mathrm{Mg}(s)+{\mathrm{Cl}}_2(g)\to {\mathrm{MgCl}}_2(s)$ Initially, we have the oxidation states as follows: - Magnesium (Mg) in its metallic state (s) has an oxidation state of 0. - Each chlorine (Cl) atom has an oxidation state of 0 since it is in its diatomic gaseous state ${\mathrm{Cl}}_2(g)$. In the product, magnesium chloride (MgCl2): - Magnesium (Mg) has an oxidation state of +2 since it typically loses two electrons to achieve the stable electron configuration of its nearest noble gas. - Each chlorine (Cl) atom has an oxidation state of -1 since it typically gains one electron to achieve the stable electron configuration of its nearest noble gas.

Determine the electron transfer during the reaction

Now that we know the initial and final oxidation states of the elements involved, we can determine how many electrons are transferred. - Magnesium (Mg) loses 2 electrons to go from an oxidation state of 0 to +2: ${\mathrm{Mg}}^0\to {\mathrm{Mg}}^{2+}+2e^{-}$ - Each chlorine (Cl) atom gains 1 electron to go from an oxidation state of 0 to -1. Since there are two chlorine atoms involved, the process happens twice: $2\times ({\mathrm{Cl}}^0+e^{-}\to {\mathrm{Cl}}^{-})$

Illustrate the electron transfer

Combining the processes above, we can illustrate the electron transfer for the reaction as follows: 1. Magnesium loses 2 electrons: ${\mathrm{Mg}}^0\to {\mathrm{Mg}}^{2+}+2e^{-}$ 2. Each chlorine gains 1 electron, (happening twice for a total of 2 electrons gained): $2\times ({\mathrm{Cl}}^0+e^{-}\to {\mathrm{Cl}}^{-})$ 3. The magnesium ion and chloride ions combine to form magnesium chloride: ${\mathrm{Mg}}^{2+}+2{\mathrm{Cl}}^{-}\to {\mathrm{MgCl}}_2$ The overall redox reaction can be written as: ${\mathrm{Mg}}^0+2{\mathrm{Cl}}^0\to {\mathrm{Mg}}^{2+}+2{\mathrm{Cl}}^{-}\to {\mathrm{MgCl}}_2$ Here, we can see that magnesium loses 2 electrons (oxidation), and each chlorine atom gains 1 electron (reduction). The electrons are thus transferred from magnesium to chlorine during the redox reaction.

Key concepts

Oxidation States

Understanding oxidation states is key to grasping the principles of redox reactions. Oxidation states, also known as oxidation numbers, indicate the degree of oxidation of an atom within a compound. It's a theoretical charge that an atom would have if electrons were completely transferred, not shared, during bond formation.

Let's take the given exercise as an example. Elemental magnesium \textbf{(Mg)} is in its pure form, so it has an oxidation state of 0. This indicates that it is neither gaining nor losing electrons. Chlorine gas \textbf{(Cl${}_2$)} also has an oxidation state of 0, as it's composed of two chlorine atoms sharing electrons equally.

However, in the product magnesium chloride (\textbf{MgCl${}_2$}), the magnesium atom has an oxidation state of +2, showing it has lost two electrons. Conversely, each chlorine atom has an oxidation state of -1, indicating that each has gained one electron. When interpreting reactions, paying attention to these changes in oxidation states allows us to track electron flow and identify which species are oxidized and which are reduced.

Electron Transfer

In a redox reaction, electron transfer is the movement of electrons from one atom to another, changing the atoms' oxidation states. This transfer of electrons is what fundamentally defines oxidation and reduction processes - collectively known as redox reactions.

For instance, during the reaction \textbf{Mg}(s) + \textbf{Cl}${}_2$(g) \rightarrow \textbf{MgCl}${}_2$(s), the magnesium atom (\textbf{Mg}) is oxidized as it loses two electrons and the chlorine atoms (\textbf{Cl}) are reduced since they each gain an electron. It is important to note that the total number of electrons lost is equal to the number gained, maintaining charge balance. One can visualize electron transfer in equations by writing half-reactions that represent the oxidation and the reduction separately.

In our case, the half-reactions would be: \textbf{Mg} \rightarrow \textbf{Mg}${}^{2+}$ + 2\textbf{e}${}^{-}$ (oxidation) and \textbf{Cl} + \textbf{e}${}^{-}$ \rightarrow \textbf{Cl}${}^{-}$ (reduction). This conceptualization helps clarify the individual changes occurring within this chemical interaction.

Chemical Reactions

Chemical reactions involve the transformation of reactants into products through the breaking and formation of bonds. Redox reactions are a specific type of chemical reaction characterized by the exchange of electrons between participating species.

In the transformation of magnesium and chlorine gas to magnesium chloride, we witness a chemical reaction where elemental substances combine to form a compound with entirely different properties. We can better understand this by dividing the redox process into two half-reactions that focus on oxidization and reduction, as described in the step-by-step solution. This approach simplifies the complex electron exchange into more manageable parts, which enhances comprehension of the complete reaction.

The disciplinal grasp of chemical reactions, especially redox reactions, spirals beyond rote memorization. Recognizing the signs of electron transfer and changes in oxidation states in various reactions underpins a solid foundation in chemistry, facilitating the interpretation and prediction of chemical behavior.