6. Chemical Equilibrium

Chemical Thermodynamics: Enthalpy

6. Chemical Equilibrium

Chemical Thermodynamics: Enthalpy: Videos & Practice Problems

Topic summary

Enthalpy, denoted as ΔH, measures heat transfer in a system at constant pressure. Exothermic processes release heat, resulting in a negative ΔH, while endothermic processes absorb heat, leading to a positive ΔH. For example, condensation and freezing are exothermic, while melting and vaporization are endothermic. Energy diagrams illustrate these changes, showing reactants and products at different energy levels. The change in enthalpy can be calculated using the equation: $Δ H = Products - Reactants$ .

Enthalpy represents the amount of kinetic energy in the form of heat between a system and its surroundings.

Enthalpy

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Enthalpy

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Enthalpy Video Summary

Enthalpy is a crucial concept in thermodynamics, representing the total kinetic energy in the form of heat exchanged between a system and its surroundings at constant pressure. The standard enthalpy of formation, denoted as ΔH_f^°, quantifies the heat change that occurs when reactants in their standard states transform into products. Understanding enthalpy involves distinguishing between exothermic and endothermic processes.

In exothermic processes, heat is released, resulting in a negative value for heat (q) and enthalpy (ΔH). For example, the reaction of three moles of hydrogen gas with one mole of nitrogen gas to produce two moles of ammonia gas illustrates this concept. Initially, reactants possess weak bonds, which, upon reaction, form stronger bonds in the products. As gas molecules lose heat, they slow down and can condense into a liquid or solid, releasing heat to the surroundings. This is why a container holding such a reaction feels warm to the touch. Common phase changes associated with exothermic reactions include condensation (gas to liquid), freezing (liquid to solid), and deposition (gas to solid).

Conversely, endothermic processes absorb heat, leading to positive values for both q and ΔH. An example is the dissociation of hydrogen bromide (HBr) into hydrogen ions (H⁺) and bromide ions (Br^-). In these reactions, strong lattice bonds require external energy to break, resulting in weaker bonds in the products. For instance, heating an ice cube allows it to absorb energy, transitioning from solid to liquid (melting or fusion) and potentially to gas (vaporization). The container in this case feels cold as the ice absorbs heat from its surroundings.

Energy diagrams effectively illustrate these concepts. In an exothermic reaction, the energy of the products is lower than that of the reactants, indicating a release of energy (ΔH is negative). In contrast, an endothermic reaction shows products at a higher energy state than reactants, reflecting energy absorption (ΔH is positive). There are also cases where reactants and products are at the same energy level, resulting in a thermoneutral reaction where ΔH equals zero.

The change in enthalpy can be summarized by the equation: ΔH = Σ(ΔH_f of products) - Σ(ΔH_f of reactants). This relationship emphasizes the energy transfer involved in chemical reactions, highlighting the significance of enthalpy in understanding thermal processes under constant pressure.

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What is the difference between exothermic and endothermic processes in terms of enthalpy?

Exothermic processes release heat to the surroundings, resulting in a negative change in enthalpy (ΔH < 0). This means the products are at a lower energy state than the reactants. Examples include condensation and freezing. Endothermic processes, on the other hand, absorb heat from the surroundings, leading to a positive change in enthalpy (ΔH > 0). This indicates that the products are at a higher energy state than the reactants. Examples include melting and vaporization. The sign of ΔH helps determine whether a process is exothermic or endothermic.

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How do you calculate the change in enthalpy (ΔH) for a chemical reaction?

The change in enthalpy (ΔH) for a chemical reaction can be calculated using the equation:

$Δ H = \frac{\sum H (products)}{- \sum H (reactants)}$

This equation means you sum the enthalpies of formation of all the products and subtract the sum of the enthalpies of formation of all the reactants. The result will give you the overall change in enthalpy for the reaction.

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What is standard enthalpy of formation (ΔH_f°) and how is it used?

The standard enthalpy of formation (ΔH_f°) is the change in enthalpy when one mole of a compound is formed from its elements in their standard states under standard conditions (298 K and 1 atm). It is used to calculate the enthalpy changes in chemical reactions by summing the standard enthalpies of formation of the products and subtracting the sum of the standard enthalpies of formation of the reactants. This helps in determining whether a reaction is exothermic or endothermic.

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What are the phase changes associated with exothermic and endothermic processes?

Exothermic phase changes release heat and include:

Condensation (gas to liquid)
Freezing (liquid to solid)
Deposition (gas to solid)

Endothermic phase changes absorb heat and include:

Melting (solid to liquid), also known as fusion
Vaporization (liquid to gas)
Sublimation (solid to gas)

These phase changes illustrate how heat transfer affects the state of matter.

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How do energy diagrams illustrate exothermic and endothermic reactions?

Energy diagrams show the energy levels of reactants and products. In an exothermic reaction, the reactants start at a higher energy level, and the products are at a lower energy level, indicating that energy is released (ΔH < 0). In an endothermic reaction, the reactants start at a lower energy level, and the products are at a higher energy level, indicating that energy is absorbed (ΔH > 0). These diagrams help visualize the energy changes during a reaction.

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