Problem 12 - Chemistry for Allied Health

Question

Balance the following equations and classify the following reactions as combination, decomposition, single replacement, double replacement, or combustion. For each, indicate if it is a redox reaction or not. a. \(\mathrm{Na}+\mathrm{Cl}_{2} \rightarrow \mathrm{NaCl}\) b. \(\mathrm{Na}_{3} \mathrm{PO}_{4}+\mathrm{KOH} \rightarrow \mathrm{NaOH}+\mathrm{K}_{3} \mathrm{PO}_{4}\) c. \(\mathrm{P}_{4}+\mathrm{O}_{2} \rightarrow \mathrm{P}_{2} \mathrm{O}_{3}\) d. \(\mathrm{N}_{2}+\mathrm{H}_{2} \rightarrow \mathrm{NH}_{3}\) e. \(\mathrm{Al}+\mathrm{HCl} \rightarrow \mathrm{H}_{2}+\mathrm{AlCl}_{3}\) f. \(\mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\) g. \(\mathrm{NH}_{3}+\mathrm{CuO} \rightarrow \mathrm{Cu}+\mathrm{N}_{2}+\mathrm{H}_{2} \mathrm{O}\) h. \(\mathrm{NH}_{4} \mathrm{NO}_{3} \rightarrow \mathrm{N}_{2} \mathrm{O}+\mathrm{H}_{2} \mathrm{O}\)

Step-by-Step Solution

Answer

Balanced equations are provided; reactions: a) combination, redox; b) double replacement, non-redox; c) combination, non-redox; d) combination, redox; e) single replacement, redox; f) decomposition, redox; g) redox; h) decomposition, redox.

Step 1: Balance Equation a

The equation is \( \mathrm{Na} + \mathrm{Cl}_2 \rightarrow \mathrm{NaCl} \). To balance, note there are 2 Cl atoms in \( \mathrm{Cl}_2 \), so you need 2 \( \mathrm{NaCl} \). The balanced equation is:\[ 2 \mathrm{Na} + \mathrm{Cl}_2 \rightarrow 2 \mathrm{NaCl} \]This is a combination reaction and a redox reaction because Na goes from 0 to +1 and Cl goes from 0 to -1.

Step 2: Balance Equation b

The equation is \( \mathrm{Na}_3 \mathrm{PO}_4 + \mathrm{KOH} \rightarrow \mathrm{NaOH} + \mathrm{K}_3 \mathrm{PO}_4 \). Balance the sodium atoms first; you need 3 \( \mathrm{NaOH} \) for 3 Na in \( \mathrm{Na}_3\mathrm{PO}_4 \). Balance potassium by ensuring there are 3 \( \mathrm{KOH} \) for 3 K in \( \mathrm{K}_3\mathrm{PO}_4 \). Balance as follows:\[ \mathrm{Na}_3 \mathrm{PO}_4 + 3 \mathrm{KOH} \rightarrow 3 \mathrm{NaOH} + \mathrm{K}_3 \mathrm{PO}_4 \]This is a double replacement reaction and not a redox reaction.

Step 3: Balance Equation c

The equation is \( \mathrm{P}_4 + \mathrm{O}_2 \rightarrow \mathrm{P}_2 \mathrm{O}_3 \). Balance phosphorus by using 2 \( \mathrm{P}_2 \mathrm{O}_3 \), then match the oxygen by using 3 \( \mathrm{O}_2 \). The balanced equation is:\[ \mathrm{P}_4 + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{P}_2 \mathrm{O}_3 \]This is a combination reaction and not a redox reaction.

Step 4: Balance Equation d

The equation is \( \mathrm{N}_2 + \mathrm{H}_2 \rightarrow \mathrm{NH}_3 \). Balance hydrogen using 6 \( \mathrm{H}_2 \) and nitrogen using 2 \( \mathrm{NH}_3 \). The balanced equation is:\[ \mathrm{N}_2 + 3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3 \]This is a combination reaction and a redox reaction because H goes from 0 to +1 and N goes from 0 to -3.

Step 5: Balance Equation e

The equation is \( \mathrm{Al} + \mathrm{HCl} \rightarrow \mathrm{H}_2 + \mathrm{AlCl}_3 \). Start by balancing Cl by ensuring 3 \( \mathrm{HCl} \) for \( \mathrm{AlCl}_3 \). Balance H by adding \( \mathrm{H}_2 \), balance using:\[ 2 \mathrm{Al} + 6 \mathrm{HCl} \rightarrow 3 \mathrm{H}_2 + 2 \mathrm{AlCl}_3 \]This is a single replacement reaction and a redox reaction as Al is oxidized and H+ reduced.

Step 6: Balance Equation f

The equation is \( \mathrm{H}_2 \mathrm{O}_2 \rightarrow \mathrm{H}_2 \mathrm{O} + \mathrm{O}_2 \). Balance oxygen by making the number of oxygens even ( multiplying waters by 2):\[ 2 \mathrm{H}_2 \mathrm{O}_2 \rightarrow 2 \mathrm{H}_2 \mathrm{O} + \mathrm{O}_2 \]This is a decomposition reaction and a redox reaction with O being both reduced and oxidized.

Step 7: Balance Equation g

The equation is \( \mathrm{NH}_3 + \mathrm{CuO} \rightarrow \mathrm{Cu} + \mathrm{N}_2 + \mathrm{H}_2 \mathrm{O} \). Balance nitrogen first, then systematic balance of H and O:\[ 2 \mathrm{NH}_3 + 3 \mathrm{CuO} \rightarrow 3 \mathrm{Cu} + \mathrm{N}_2 + 3 \mathrm{H}_2 \mathrm{O} \]This is a redox reaction (NH3 is oxidized, CuO is reduced).

Step 8: Balance Equation h

The equation is \( \mathrm{NH}_4 \mathrm{NO}_3 \rightarrow \mathrm{N}_2 \mathrm{O} + \mathrm{H}_2 \mathrm{O} \). Balance using stoichiometric coefficients:\[ \mathrm{NH}_4 \mathrm{NO}_3 \rightarrow \mathrm{N}_2 \mathrm{O} + 2 \mathrm{H}_2 \mathrm{O} \]This is a decomposition reaction that involves an internal redox process.

Key Concepts

Redox Reactions

Redox reactions are a fundamental concept in chemistry. They involve the transfer of electrons between substances. In a redox reaction, one substance is oxidized, meaning it loses electrons, while another is reduced, meaning it gains electrons. This electron exchange results in a change in oxidation states of the elements involved.
For example, in the balanced reaction of \(2 \mathrm{Na} + \mathrm{Cl}_2 \rightarrow 2 \mathrm{NaCl} \), sodium \(\mathrm{Na} \) goes from an oxidation state of 0 to +1 (oxidized), and chlorine \(\mathrm{Cl}_2\) goes from 0 to -1 (reduced). Such changes in oxidation states confirm the occurrence of a redox reaction.

Redox reactions are crucial in numerous chemical processes, including energy generation in batteries, metabolic pathways in living organisms, and various industrial processes. Understanding how to identify and balance these reactions is essential for recognizing their role in chemistry and everyday life.

Reaction Classification

Classifying reactions helps in understanding the processes involved in chemical changes. There are several distinct types of chemical reactions based on their characteristics and changes involved.

• Combination Reactions: Two or more substances combine to form a single product. Example: \(\mathrm{N}_2 + 3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3\).
• Decomposition Reactions: A single compound breaks down into two or more products. Example: \(2 \mathrm{H}_2 \mathrm{O}_2 \rightarrow 2 \mathrm{H}_2 \mathrm{O} + \mathrm{O}_2\).
• Single Replacement Reactions: An element replaces another element in a compound. Example: \(2 \mathrm{Al} + 6 \mathrm{HCl} \rightarrow 3 \mathrm{H}_2 + 2 \mathrm{AlCl}_3\).
• Double Replacement Reactions: Components of two compounds exchange places. Example: \(\mathrm{Na}_3 \mathrm{PO}_4 + 3 \mathrm{KOH} \rightarrow 3 \mathrm{NaOH} + \mathrm{K}_3 \mathrm{PO}_4\).
• Combustion Reactions: Rapid reactions with oxygen that release energy as light or heat.

Knowing the type of reaction helps in predicting the products and understanding the underlying chemical processes better.

Types of Chemical Reactions

There are diverse types of chemical reactions that occur based on different criteria, including energy changes, atoms' movements, or the nature of the reactants.
Each type of reaction serves a specific purpose and function. Here’s more detail on them:

• Combination Reactions: Common in synthesis processes where multiple reactants form a complex product.
• Decomposition Reactions: Essential for breaking down compounds, often requiring an energy input like heat or electricity.
• Single Replacement Reactions: These are driven by the reactivity of the involved elements and are commonly seen in metallurgy and electrochemical cells.
• Double Replacement Reactions: Found where ion exchange occurs, such as in neutralization or precipitation reactions.

Understanding the type of reaction can also imply the energy requirements or releases associated, which is key for applications in energy generation and storage.

Chemical Equation Balancing Methods

Balancing chemical equations is a crucial skill in chemistry that ensures the law of conservation of mass is respected. It involves making sure that the number of each type of atom is the same on both sides of the equation.
Here are some tips for balancing:

• Begin by balancing atoms of elements that appear in only one reactant and product.
• Next, balance atoms that appear in multiple reactants or products using the least common multiple method.
• Leave hydrogen and oxygen for last, as they are often found in multiple compounds.
• Adjust coefficients, not subscripts, to balance the equation.

For example, to balance the equation \(\mathrm{N}_2 + 3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3\), you start by determining that nitrogen is balanced, then adjust hydrogen accordingly.

Mastering these methods is essential in predicting reaction outcomes, optimizing reaction conditions, and understanding reaction mechanisms accurately.