Question

(a) Using average bond enthalpies, predict which of the following reactions will be most exothermic: (i) \(\mathrm{C}(\mathrm{g})+2 \mathrm{~F}_{2}(\mathrm{~g}) \longrightarrow \mathrm{CF}_{4}(\mathrm{~g})\) (ii) \(\mathrm{CO}(\mathrm{g})+3 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+\mathrm{OF}_{2}(g)\) (iii) \(\mathrm{CO}_{2}(g)+4 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+2 \mathrm{OF}_{2}(g)\) (b) Explain the trend, if any, that exists between reaction exothermicity and the extent to which the carbon atom is bonded to oxygen.

Short answer

The most exothermic reaction would be the one with the most negative ∆H value. By calculating the ∆H for each reaction using bond enthalpies, we can compare the values and find the most exothermic reaction. When examining the carbon-oxygen bond strength, we can observe that the stronger the bond, the more energy is required to break those bonds, which can affect the reaction's exothermicity. Analyzing the enthalpy changes and carbon-oxygen bond strength can help us understand the trend between reaction exothermicity and the extent of carbon-oxygen bonding.

Step-by-step solution

Understanding ∆H and bond enthalpy relationships

We must understand that the enthalpy change of a reaction (∆H) can be calculated using the bond enthalpies of the participating species. ∆H is calculated as the difference between the bond enthalpies of the reactants and the products: \[∆H = ∑(bond\ enthalpies\ of\ reactants) - ∑(bond\ enthalpies\ of\ products)\] We'll now apply this formula to each of the given reactions.

Calculating ∆H for Reaction (i)

For Reaction (i), we have: \[C(g) + 2F_2(g) \longrightarrow CF_4(g)\] In this reaction, we have C with all single bonds to the F atoms in CF4. We must consider both the carbon-fluorine (C-F) bonds and the fluorine-fluorine (F-F) bonds in F2 when calculating the bond enthalpies. In this case: \[∆H_i = (C_{bond\ enthalpy} + 2 \times F_{2\ bond\ enthalpy}) - 4 \times (C-F\ bond\ enthalpy)\]

Calculating ∆H for Reaction (ii)

For Reaction (ii), we have: \[CO(g) + 3F_2(g) \longrightarrow CF_4(g) + OF_2(g)\] In this reaction, we must account for the carbon-oxygen (C=O) bonds in CO, the F-F bonds in F2, and the carbon-fluorine (C-F) bonds in CF4, and the oxygen-fluorine (O-F) bonds in OF2 when calculating bond enthalpies. In this case: \[∆H_{ii} = (C=O\ bond\ enthalpy + 3 \times F_{2\ bond\ enthalpy}) - (4 \times (C-F\ bond\ enthalpy) + 2 \times (O-F\ bond\ enthalpy))\]

Calculating ∆H for Reaction (iii)

For Reaction (iii), we have: \[CO_2(g) + 4F_2(g) \longrightarrow CF_4(g)+ 2OF_2(g)\] In this reaction, we must account for the carbon-oxygen (C=O) bonds in CO2, the F-F bonds in F2, the carbon-fluorine (C-F) bonds in CF4, and the oxygen-fluorine (O-F) bonds in OF2 when calculating bond enthalpies. In this case: \[∆H_{iii} = (2 \times C=O\ bond\ enthalpy + 4 \times F_{2\ bond\ enthalpy}) - (4 \times (C-F\ bond\ enthalpy) + 4 \times (O-F\ bond\ enthalpy))\]

Comparing ∆H values

Once we have calculated the enthalpy changes (∆H) for each reaction, we compare them to find out which reaction is the most exothermic. The reaction having the most negative value of ∆H will be the most exothermic.

Analyzing the relationship between reaction exothermicity and carbon-oxygen bonding

To explain the trend between reaction exothermicity and the extent to which the carbon atom is bonded to oxygen, we have to observe the differences in bond strengths between carbon and oxygen in the various reactants and products. More specifically, we need to focus on the dissociation energy of bonds involving carbon and oxygen and compare it with other reactions. Generally, the stronger the bond between carbon and oxygen atoms, the more energy is required to break those bonds, which can also affect the reaction's exothermicity. To summarize, by comparing the enthalpy changes for each reaction and analyzing the carbon-oxygen bond strength in the involved molecules, we can predict which reaction is the most exothermic and understand the trend between the reaction exothermicity and the carbon-oxygen bonding extent.

Key concepts

Enthalpy Change

In chemistry, enthalpy change, denoted as \( \Delta H \), is a measure of the heat absorbed or released during a chemical reaction at constant pressure. It's a central concept in thermodynamics which tells us a lot about the energy dynamics of a chemical process. To grasp this idea, let's compare it to your everyday experience: think of it like the ‘financial cost’ of a reaction – some reactions ‘pay out’ energy (exothermic), while others 'cost' energy (endothermic).

Understanding enthalpy change is crucial for predicting whether a reaction will release heat to the surroundings, as seen in exothermic reactions, or absorb heat, characterizing endothermic reactions. In the context of the exercise, we used the formula \( \Delta H = \sum (\text{bond enthalpies of reactants}) - \sum (\text{bond enthalpies of products}) \) to calculate the enthalpy changes of different reactions involving carbon and fluorine. This calculation is akin to balancing a checkbook, where you ensure that the energy put in and taken out aligns correctly.

• The sum of the bond enthalpies of reactants is the 'energy input' required to break existing bonds.
• The sum of the bond enthalpies of products is the 'energy output' released when new bonds are formed.

In a chemical reaction, breaking bonds always requires energy (endothermic process), while forming new bonds releases energy (exothermic process). Therefore, if the 'energy output' is greater than the 'energy input', the overall \( \Delta H \) will be negative, indicating an exothermic reaction. Conversely, if more energy is needed to break the bonds than is released when new ones are formed, the reaction is endothermic, with a positive \( \Delta H \).
The calculations from the exercise provide a systematic approach to determine the enthalpy change, ultimately helping to predict how much energy a reaction will release or consume.

Exothermic Reactions

Heat, in the world of chemistry, isn't just something that warms our hands; it's a critical indicator of what's happening at the molecular level during a reaction. Exothermic reactions are the celebrities of chemical processes - they are the ones giving off heat, making their presence known with warmth and sometimes even light.

Envision a campfire: as the wood burns, it releases heat into the cool night, which is an example of an exothermic reaction taking place. In these types of reactions, the end products are more stable than the reactants, translating to a release of energy. This is similar to relaxation—when things settle down, they tend to give off the excess energy they no longer need.

Characteristics of Exothermic Reactions

• Release of heat to the surroundings.
• Typically observable as warming or heating.
• Results in a decrease in the energy content of the system, hence a negative \( \Delta H \).
• Favorable in terms of energy, often spontaneous.

What makes a reaction exothermic? It's all about the bonds. Stronger bonds being formed in the product than those broken in the reactants will lead to an exothermic reaction because forming a bond releases more energy than is required to break a bond.
In our exercise, by analyzing the enthalpy changes (\( \Delta H \)) using bond enthalpies, we determine which one of the presented reactions will be the most exothermic—this is the reaction that effectively 'pays out' the most energy, revealing itself as the reaction that heats up its surroundings the most.

Chemical Bonding

The 'glue' that holds the atoms in molecules together is what we refer to as chemical bonding. At its core, chemical bonding is the physical phenomenon of chemical substances being held together by attraction of atoms to each other through sharing, as well as exchanging, of electrons.

Think of atoms like individuals at a dance. Some prefer to dance alone (noble gases), while others are eager to form partnerships to become more stable (such as in covalent or ionic bonds). In our exercise, the dance involves carbon (C), oxygen (O), and fluorine (F), each 'dancing' to the tune of chemical stability by forming and breaking bonds.

Types of Chemical Bonds

• Covalent Bonds: Atoms share electron pairs, as seen in molecules like \( CF_4 \).
• Ionic Bonds: Atoms transfer electrons to achieve stability, typically between metals and non-metals.
• Metallic Bonds: Free-flowing electrons in metals lead to a sea of electrons scenario, giving metals their characteristic properties.

Bonding dictates the properties of substances and the energy changes occurring during a reaction. For instance, in our exercise, we see how the different strengths of covalent bonds between carbon and either oxygen or fluorine influence the overall bond enthalpies and subsequently the reaction's enthalpy change. The dance of atoms, their bonding patterns, directly relates to how much energy will be consumed or released in a reaction, impacting whether a reaction will be exothermal or endothermal.
Identifying these bonding interactions helps us understand how molecules will react with one another, predict the energies involved, and determine the stability of the formed compounds. By observing the carbon-fluorine and carbon-oxygen bonds in various scenarios, students can appreciate how these bonds impact the reactivity and the heat released during reactions.