Question

Although the metals of Group 2 of the periodic table are not nearly as reactive as those of Group \(1,\) many of the Group 2 metals will combine with common nonmetals, especially at elevated temperatures. Write balanced chemical equations for the reactions of \(\mathrm{Mg}, \mathrm{Ca}, \mathrm{Sr}\) and Ba with \(\mathrm{Cl}_{2}, \mathrm{Br}_{2},\) and \(\mathrm{O}_{2}\).

Short answer

The balanced chemical equations for the reactions of Mg, Ca, Sr, and Ba with Cl₂, Br₂, and O₂ are as follows: Reactions with Cl₂: 1. Mg + Cl₂ → MgCl₂ 2. Ca + Cl₂ → CaCl₂ 3. Sr + Cl₂ → SrCl₂ 4. Ba + Cl₂ → BaCl₂ Reactions with Br₂: 5. Mg + Br₂ → MgBr₂ 6. Ca + Br₂ → CaBr₂ 7. Sr + Br₂ → SrBr₂ 8. Ba + Br₂ → BaBr₂ Reactions with O₂: 9. 2Mg + O₂ → 2MgO 10. 2Ca + O₂ → 2CaO 11. 2Sr + O₂ → 2SrO 12. 2Ba + O₂ → 2BaO

Step-by-step solution

Find the charges of the ions

For Group 2 metals, such as Mg, Ca, Sr, and Ba, they have an oxidation state of +2 due to the presence of 2 valence electrons, which are lost during the reaction. For nonmetals, we can look at their position in the periodic table. Cl (Group 17) and Br (Group 17) will each have a charge of -1. Oxygen (Group 16) will have a charge of -2.

Write the general form of the reactions

Using the oxidation states from Step 1, we can write the general form of the reactions: - Metal + Cl₂ → Metal(II) Chloride - Metal + Br₂ → Metal(II) Bromide - Metal + O₂ → Metal(II) Oxide

Write the balanced chemical equations for each reaction

Reactions with Cl₂: 1. Mg + Cl₂ → MgCl₂ 2. Ca + Cl₂ → CaCl₂ 3. Sr + Cl₂ → SrCl₂ 4. Ba + Cl₂ → BaCl₂ Reactions with Br₂: 5. Mg + Br₂ → MgBr₂ 6. Ca + Br₂ → CaBr₂ 7. Sr + Br₂ → SrBr₂ 8. Ba + Br₂ → BaBr₂ Reactions with O₂: 9. 2Mg + O₂ → 2MgO 10. 2Ca + O₂ → 2CaO 11. 2Sr + O₂ → 2SrO 12. 2Ba + O₂ → 2BaO These balanced chemical equations represent the reactions of Mg, Ca, Sr, and Ba with Cl₂, Br₂, and O₂ respectively.

Key concepts

Group 2 Metals Reactivity

When studying Chemistry, understanding the reactivity of elements is crucial. Group 2 metals, also known as the alkaline earth metals, include magnesium (Mg), calcium (Ca), strontium (Sr), and barium (Ba). These metals are known for being less reactive than their Group 1 counterparts, but they still participate in various chemical reactions, particularly with nonmetals.

What makes these metals reactive? It's primarily due to the two valence electrons they possess. These electrons are relatively easy to lose, making these metals good reducing agents. For instance, in reactions where these metals are exposed to halogens like chlorine (Cl₂) or bromine (Br₂), or even to a simple diatomic molecule like oxygen (O₂), they readily give up their electrons.

Such reactions are often more vigorous at elevated temperatures, enabling the metal atoms to overcome activation energy barriers and react to form ionic compounds. The reactivity of Group 2 metals increases as you move down the periodic table; thus, barium will react more readily with a given nonmetal compared to magnesium.

Oxidation States

Delving into chemical reactions means that you'll often encounter the term 'oxidation states.' This concept is pivotal in understanding how elements combine to form compounds. The oxidation state, sometimes called oxidation number, is an indicator of the degree of oxidation (loss of electrons) of an atom in a chemical compound.

In the context of Group 2 metals — Mg, Ca, Sr, and Ba, they always have an oxidation state of +2. This is because they lose two electrons to achieve a full outer shell, which is the most stable arrangement of electrons. In a chemical equation, this is represented by a Roman numeral after the metal: for example, Mg^II indicates magnesium in a +2 oxidation state.

Understanding oxidation states helps in predicting how elements will interact. For nonmetals, we consider their group number to determine their most common oxidation state. For instance, halogens like Cl and Br typically have an oxidation state of -1, while oxygen usually has an oxidation state of -2. This information is fundamental when balancing chemical equations to ensure that the total charge is the same on both sides of the reaction.

Chemical Reactions of Metals with Nonmetals

The reactions between metals and nonmetals are some of the most fundamental and widespread chemical processes. These reactions often result in the formation of ionic compounds, which are substances composed of oppositely charged ions. In our scenario with Group 2 metals such as Mg, Ca, Sr, and Ba, when they react with nonmetals like chlorine (Cl₂), bromine (Br₂), or oxygen (O₂), the metals lose electrons to become positively charged ions, while the nonmetals gain electrons to become negatively charged ions.

To write chemical equations for these reactions, we use the oxidation states to determine the formula of the resultant compounds. For instance:

• Mg and Cl₂ react to form MgCl₂, where magnesium loses two electrons to chlorine.
• Similarly, Ca and O₂ react to form CaO, with calcium giving up two electrons to oxygen.

It is essential to note that in reactions with oxygen, the metals form oxides with the general formula MetalO, and two metal atoms will react with one oxygen molecule (O₂). This is due to the -2 oxidation state of oxygen, requiring two metal atoms for balance.

Writing balanced chemical equations ensures that the number of atoms for each element is the same on both sides of the equation, reflecting the law of conservation of mass. For example:

• 2Mg + O₂ → 2MgO
• 2Ca + O₂ → 2CaO

Each balanced equation exhibits a clear exchange of electrons between the metals and nonmetals, illustrating the reducing nature of Group 2 metals and the oxidizing nature of the nonmetals involved.