Question

Without doing any calculations, predict the sign of \(\Delta H\) for each of the following reactions: $$\begin{array}{l}{\text { (a) } \mathrm{NaCl}(s) \longrightarrow \mathrm{Na}^{+}(g)+\mathrm{Cl}^{-}(\mathrm{g})} \\ {\text { (b) } 2 \mathrm{H}(g) \longrightarrow \mathrm{H}_{2}(g)} \\ {\text { (c) } \mathrm{Na}(g) \longrightarrow \mathrm{Na}^{+}(g)+\mathrm{e}^{-}} \\ {\text { (d) } \mathrm{I}_{2}(s) \longrightarrow \mathrm{I}_{2}(l)}\end{array}$$

Short answer

For the given reactions, the signs of \(\Delta H\) are as follows: - Reaction (a): Positive \(\Delta H\) - endothermic - Reaction (b): Negative \(\Delta H\) - exothermic - Reaction (c): Positive \(\Delta H\) - endothermic - Reaction (d): Positive \(\Delta H\) - endothermic

Step-by-step solution

Analyze the reaction (a) NaCl(s) -> Na+(g) + Cl-(g)

In this reaction, a solid ionic compound (NaCl) is being dissociated into its gas-phase ions (Na+ and Cl-). This process requires energy to break the ionic bonds between the solid lattice. Therefore, this reaction will absorb heat from the surroundings, making it an endothermic reaction.

Predict the sign of ΔH for reaction (a)

Since reaction (a) is endothermic, the change in enthalpy, \(\Delta H\), is positive.

Analyze the reaction (b) 2 H(g) -> H2(g)

In this reaction, two gaseous hydrogen atoms are combining to form a gaseous hydrogen molecule (H2). The process involves the formation of a bond between two hydrogen atoms, which releases energy. Therefore, this reaction is exothermic.

Predict the sign of ΔH for reaction (b)

Since reaction (b) is exothermic, the change in enthalpy, \(\Delta H\), is negative.

Analyze the reaction (c) Na(g) -> Na+(g) + e-

In this reaction, a neutral gaseous sodium atom is being ionized to lose an electron, forming a gaseous sodium ion (Na+). Energy is required to overcome the electrostatic attraction between the electron and the positively charged nucleus. Therefore, this reaction is endothermic.

Predict the sign of ΔH for reaction (c)

Since reaction (c) is endothermic, the change in enthalpy, \(\Delta H\), is positive.

Analyze the reaction (d) I2(s) -> I2(l)

In this reaction, solid iodine (I2) is being converted into liquid iodine. When a substance changes from a solid to a liquid phase, energy is required to overcome the forces keeping the particles organized in the solid phase. This is called the heat of fusion. Therefore, this reaction is endothermic.

Predict the sign of ΔH for reaction (d)

Since reaction (d) is endothermic, the change in enthalpy, \(\Delta H\), is positive. In summary: - Reaction (a): Positive \(\Delta H\) - Reaction (b): Negative \(\Delta H\) - Reaction (c): Positive \(\Delta H\) - Reaction (d): Positive \(\Delta H\)

Key concepts

Endothermic Reactions

Endothermic reactions are a fascinating aspect of chemistry where the reaction absorbs heat from its surroundings. Imagine it like a sponge soaking up water – instead of releasing heat, these reactions require heat to proceed.

In endothermic reactions, the products have higher energy than the reactants because energy is absorbed. This leads to a positive change in enthalpy, or \(\Delta H\).

For example, during the dissociation of sodium chloride (NaCl) into its ionic components, as seen in reaction (a) \(\text{NaCl(s) \rightarrow Na}^{+}\text{(g) + Cl}^{-}\text{(g)}\), energy is required to break the strong ionic bonds between the atoms.

Other characteristics of endothermic reactions include:

• The surrounding temperature drops as energy is taken in.
• These reactions typically occur when bonds are broken, such as melting or vaporization."

Understanding endothermic processes is crucial in fields such as thermodynamics and materials science. They explain phenomena like why ice absorbs heat to melt.

Exothermic Reactions

Exothermic reactions are the opposite of endothermic reactions, as they release energy into the surroundings, usually in the form of heat.

These reactions have products with lower energy than the reactants, resulting in a negative \(\Delta H\). Picture a warm campfire—the burning wood undergoes an exothermic process, providing warmth and light.

In reaction (b) \(\text{2 H(g) \rightarrow H}_2\text{(g)}\), energy is released when two hydrogen atoms form a bond, creating an \(\text{H}_2\) molecule. This helps illustrate how bond formation typically accompanies exothermic reactions.

Key features of exothermic reactions include:

• An increase in surrounding temperature as heat is released.
• These reactions are usually more spontaneous and may include combustion and oxidation processes.

Exothermic processes are essential for understanding everyday phenomena, such as why certain reactions feel warm or why explosions occur.

Enthalpy Change

Enthalpy change is a core concept in thermodynamics that refers to the heat change at constant pressure during a chemical reaction. It's represented by the symbol \(\Delta H\).

Whether a reaction is endothermic or exothermic profoundly impacts the sign of \(\Delta H\).

For endothermic reactions, the enthalpy change is positive because the system absorbs energy. This is evident in reactions like the breaking of bonds or phase changes that require heat input, such as the conversion of solid iodine to liquid (\(\text{I}_2\text{(s) \rightarrow I}_2\text{(l)}\) in reaction (d)).

In contrast, exothermic reactions have a negative \(\Delta H\) as they release energy, typically during bond formation.

Some important points about enthalpy change include:

• \(\Delta H\) provides insight into the heat flow of a reaction.
• It's important in determining reaction feasibility and spontaneity.
• Enthalpy changes can inform energy requirements for processes involving heating or cooling.

Understanding \(\Delta H\) helps us to comprehend energy transformations in chemical reactions and their practical applications.