Question

Write the equation for the reaction of iron metal with hydrobromic acid to form iron(III) bromide and hydrogen gas. Determine the net change in oxidation for the element that is reduced and the element that is oxidized.

Short answer

The balanced chemical equation for the reaction between iron metal (Fe) and hydrobromic acid (HBr) to form iron(III) bromide (FeBr₃) and hydrogen gas (H₂) is: 2 Fe + 6 HBr -> 2 FeBr₃ + 3 H₂. In this reaction, iron is oxidized from an oxidation state of 0 to +3, while hydrogen is reduced from an oxidation state of +1 to 0. The net change in oxidation for the reduced element (Hydrogen) is -6, and for the oxidized element (Iron) is +3.

Step-by-step solution

Write the balanced chemical equation

First, we need to write the balanced chemical equation for the reaction between iron (Fe) and hydrobromic acid (HBr) to form iron(III) bromide (FeBr₃) and hydrogen gas (H₂). The unbalanced equation is: Fe + HBr -> FeBr₃ + H₂ To balance the equation, we need to consider the stoichiometry of the reactants and products: 2 Fe + 6 HBr -> 2 FeBr₃ + 3 H₂

Identify the elements that are reduced and oxidized

Next, we will determine the oxidation states of the reactants and products to identify the elements that undergo a change in their oxidation state. The oxidation states for each element are as follows: Fe: 0 (in elemental form) H: +1 (in HBr) Br: -1 (in HBr) Fe: +3 (in FeBr₃) H: 0 (in elemental form as H₂) Iron (Fe) changes its oxidation state from 0 to +3, which indicates that it is oxidized. Hydrogen (H) changes its oxidation state from +1 to 0, which indicates that it is reduced.

Calculate the net changes in oxidation for the reduced and oxidized elements

Now, we will calculate the net change in oxidation for each element: For Iron (Fe): Initial oxidation state = 0 Final oxidation state = +3 Net change in oxidation = (+3) - (0) = +3 For Hydrogen (H): Initial oxidation state = +1 Final oxidation state = 0 Net change in oxidation = (0) - (+1) = -1 However, since there are 6 HBr molecules, the total net change in oxidation for hydrogen is -1 times 6, which is -6. Therefore, the net change in oxidation for the element that is reduced (Hydrogen) is -6, and for the element that is oxidized (Iron) is +3.

Key concepts

Oxidation State

Oxidation state, also known as oxidation number, is a fundamental concept in chemistry that describes the degree of oxidation of an atom within a compound. It is an indicative number that implies the number of electrons an atom either gains or loses when it forms a chemical compound or ion.

Getting a clear picture of oxidation states helps students understand the changes that atoms undergo during a chemical reaction. One important thing to remember is that in its elemental form, an atom always has an oxidation state of zero. For instance, in the reaction between iron and hydrobromic acid, iron (Fe) in its metallic form has an oxidation state of 0. On the other hand, within compounds, the oxidation state of each element is determined based on a set of rules, with some common ones being hydrogen is usually +1 (except in metal hydrides where it's -1), and halogens like bromine are typically -1 (except when they're more positive in compounds with oxygen or other halogens).

In the provided exercise, identifying the change in oxidation state from reactants to products allowed us to infer that iron is oxidized (its oxidation state increases from 0 to +3), and hydrogen is reduced (its oxidation state decreases from +1 to 0). These changes are essential for balancing redox reactions, which directly leads to the stoichiometry of the equation.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are chemical processes where the oxidation state of atoms changes through the transfer of electrons. In these reactions, oxidation and reduction occur simultaneously. Oxidation involves the loss of electrons, resulting in an increase in oxidation state, while reduction involves the gain of electrons, leading to a decrease in oxidation state.

The reaction between iron and hydrobromic acid to produce iron(III) bromide and hydrogen gas is a classic example of a redox process. Iron goes from an oxidation state of 0 to +3, indicating it has lost electrons – it has been oxidized. Conversely, hydrogen's oxidation state changes from +1 to 0, showing it has gained electrons – it has been reduced.

Identifying the Redox Pair

To fully grasp redox reactions, it's crucial to pinpoint the 'redox pair,' which is made up of the oxidizing and reducing agents. The substance that gets reduced is the oxidizing agent because it causes the other substance to lose electrons. Simultaneously, the substance that gets oxidized is the reducing agent as it provides electrons to the substance being reduced. This understanding not only explains the electron flow in a reaction but is also integral for predicting possible products of unknown reactions.

Stoichiometry

Stoichiometry refers to the quantitative relationship between reactants and products in a chemical reaction. It involves using the balanced chemical equation to calculate how much of each substance is consumed or produced. Stoichiometry takes into account the conservation of mass and the principle that matter is neither created nor destroyed in chemical reactions.

For students trying to balance the chemical equations of redox reactions, stoichiometry is like the roadmap that leads to the correct coefficients for each compound. In the reaction between iron (Fe) and hydrobromic acid (HBr), the stoichiometry of the equation indicates that two moles of iron react with six moles of hydrobromic acid to produce two moles of iron(III) bromide and three moles of hydrogen gas.

Conversion Factors

Balancing chemical equations requires us to use stoichiometry as a tool for conversion, changing moles into grams or liters, depending on the states of the reactants and products. It's essential to use the molar masses of the substances involved to move from the atomic scale to a more palpable, measurable scale. Understanding stoichiometry is vital in predicting the amount of each substance needed in a reaction, identifying limiting reactants, and determining the percentage yield of products, making it a central aspect of practical chemistry.