Mastering Enthalpy: Your Guide to Delta H in Chemistry
Dive into the world of enthalpy and delta H! Understand positive and negative enthalpy changes, explore equations, and learn to predict chemical reactions. Perfect for students seeking to excel in thermodynamics.

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Now Playing:Enthalpy – Example 0a
Intros
  1. Enthalpy changes
  2. Exothermic and Endothermic reactions
  3. Mean bond enthalpy and other enthalpy definitions.
Examples
  1. Recall how to represent the enthalpy change of a reaction.
    The reaction between fluorine gas, F2 and hydrogen gas, H2 is as follows. The enthalpy change of the reaction is -539 kJ/mol:

    H2(g)+F2(g)2HF(g) \mathrm{H_{2(g)} + F_{2(g)} \to 2HF_{(g)}}

    1. i) How can the enthalpy change, ΔH, of the reaction be shown in a reaction equation?
      ii) Name one way in which the enthalpy of reaction could be measured.

    2. How could this enthalpy change be shown on an energy diagram?

Introduction to kinetics
Notes

In this lesson, we will learn:

  • To understand what is meant by exothermic and endothermic reactions in terms of heat.
  • To understand how bond enthalpies are used to find total enthalpy changes in reactions.
  • To understand how enthalpy changes are represented in chemistry.
Notes:

  • Recall the two ways (mentioned in lessons on Introduction to kinetics) to classify chemical reactions in terms of enthalpy – heat energy of a system:
    • Exothermic reactions are chemical reactions which have the overall effect of releasing heat energy to the environment. In other words, the energy put in that was needed to break up the reactant bonds was less than the energy given out when the new bonds in the products formed. This is shown by the energy diagram below:

      Enthalpy

    • Endothermic reactions are chemical reactions which have the overall effect of absorbing heat energy from the environment. In other words, the energy given out when the products formed was less than what was needed to break up the bonds in the reactants. This is shown by the energy change diagram below:

      Enthalpy


    Note the language overall release/absorbing of energy. All chemical reactions both absorb energy to break reactants and release energy when forming products. The difference between exothermic and endothermic reactions is the amounts in those two steps: exothermic releases more, endothermic absorbs more.

  • Energy in those diagrams above means chemical potential energy. This the combined attractive and repulsive forces in the atom or molecule. Attractive forces decrease this potential energy and repulsive forces increase it. This is why high energy molecules are thought of as being less stable than low energy ones.

  • It's easiest to think of this in terms of bond breaking and making:
    • Ionic and covalent bonds are strong attractive forces that create a stable low energy state. To break those bonds, energy must be put in to bring the bond/molecule to a high energy unstable state and overcome the attractive forces.
      Once the bond is broken, the atoms that were making the bond are broken up in an isolated, higher energy state possessing the energy you just put in. We say the process of breaking bonds costs energy.
    • When a strong covalent bond (like a C-C bond) forms, think of what the bond is like for the atoms compared to when they are not in a bond. The atoms now have a strong attractive force between them which stabilizes them its a much lower energy state than the high energy, uncombined state before they made the bond. This surplus energy from their old higher energy states is given off to the environment as heat. We say the process of forming bonds releases heat energy to the environment.

  • The absolute enthalpy of a substance is hard to measure, and in most cases, we only want to know it so we can find the difference in enthalpy between two substances or groups of substances, like the reactants and products of a reaction. We often just want to know the enthalpy change.
    Enthalpy changes are easier to find. If you think of a reaction as breaking some bonds (in the reactants) and forming others (in the products), finding the energy spent doing that is a good place to start!
    Mean bond enthalpy is the amount of energy required to break one mole of a chemical bond (one mole of molecules of the bond being talked about). This is usually taken as an average from several similar molecules that contain the bond being investigated.
    • This is the easiest measure of a bonds strength. A high bond enthalpy means a lot of energy is required to break this type of bond (for example the NN triple bond or the CO bond).
    • This definition is about breaking bonds, so if we are looking at products and we know the mean bond enthalpy of the bonds in the products.

  • There are other important enthalpy definitions that should be known:
    The standard enthalpy change of reaction, ΔHr\Delta H^{\ominus} _{r} , is the enthalpy change measured when a reaction takes place at standard conditions with all reactants and products in their standard states.
    • Standard conditions are defined as 100 kPa and 1M concentration if a solution. 298K is not strictly in the definition but is usually used for more consistency.
    • The standard state is the state (solid/liquid/gas) that the substance is most stable and frequently found at under standard conditions, like oxygen gas and liquid water.

    The standard enthalpy of reaction is a very general term. Youll normally use a specific version of it for the type of reaction being studied:
    • The standard enthalpy of combustion, ΔHc\Delta H^{\ominus} _{c} , is the enthalpy change when one mole of a substance is completely reacted with oxygen at standard conditions.
    • The standard enthalpy of neutralization, ΔHn\Delta H^{\ominus} _{n} , is the enthalpy change when an acid and base react to form one mole of water in a neutralization reaction at standard conditions.
    • The standard enthalpy of formation, ΔHf\Delta H^{\ominus} _{f} , is the enthalpy change when one mole of substance is formed from its elements in their standard states at standard conditions.

    Notice how they all contain one mole and standard conditions? These definitions are consistent so we can easily compare other data. We know if we look in a data book for some values to use in calculation that the values are for doing a reaction with the same amount of substance (one mole) at the same set of conditions (100 kPa, 1M concentration and almost always 298 K).

  • Enthalpy change of a reaction can be found by several equations depending on what data you have:

  • Using the enthalpy of combustion:

    ΔH=ΔHc\Delta H = \Delta H^{\ominus} _{c} (reactants) - ΔHc\Delta H^{\ominus} _{c} (products)

    Using the enthalpy of formation:

    ΔH=ΔHf\Delta H = \Delta H^{\ominus} _{f} (reactants) - ΔHf\Delta H^{\ominus} _{f} (products)

    Using mean bond enthalpy values:

    ΔH=ΔH\Delta H = \sum \Delta H (bonds broken) - ΔH=ΔH\Delta H = \sum \Delta H (bonds formed)

    Dont get these equations mixed up! The order of products and reactants in each equation is very important and will be looked at more in the next two lessons.

    Think about what the sign of the enthalpy change will be from these equations.
    The sign of the enthalpy change tells you whether a reaction is exothermic or endothermic:
    • A negative ΔH \Delta H value means a reaction is exothermic, because the enthalpy of the products is less than the enthalpy of the reactants. This means that the reaction produced lower energy products from higher energy reactants, and the change in heat was released from the system to the surroundings.
    • A positive ΔH \Delta H value means a reaction is endothermic, because the enthalpy of the products is greater than the enthalpy of the reactants. This means that the reaction produced higher energy products from lower energy reactants, and this net higher energy change was energy absorbed from the surroundings.

  • Enthalpy changes for chemical reactions can be presented in a few ways:
    • In the form of an equation, for example:

    • A+B+A + B + 25kJ    \; C \, C \qquad Also written as:
      A+B  A + B \; CΔH=+ \, C \qquad \Delta H = \, + 25kJ/mol

    • Using an energy change diagram as shown below. The enthalpy change value is marked by the final difference between reactant(s) and product(s) energy.

    Enthalpy

    The diagram above is different to those at the top of the page as it includes the activation energy of the reaction.

  • There are several ways to measure the enthalpy change of a chemical reaction. Depending on what reaction you are studying, you can use:
    • Mean bond enthalpy, (Bond-enthalpy) which is using the sum of bonds broken in reactants against the sum of bonds formed in products to find the enthalpy change.
    • Hesss law (Calculating enthalpy: Hesss Law): if you dont know the enthalpy change from AA \, B \, B but you do know it from AA \, C \, C and from BB \, C \, C, you can find AA \, B \, B out indirectly.
    • Calorimetry (Calorimetry), which measures enthalpy change based on temperature change in a reactions surroundings.
Concept

Introduction to Enthalpy

Welcome to our exploration of enthalpy, a crucial concept in chemistry! Enthalpy, often denoted as H, is a measure of the total heat content of a system. When we talk about changes in enthalpy (delta H), we're discussing the heat absorbed or released during a chemical reaction or physical process. Understanding enthalpy is essential for grasping thermodynamics and predicting the direction of chemical reactions. To kick off our journey, I've prepared an introduction video that breaks down this complex topic into digestible bits. This video will help you visualize enthalpy changes and their significance in various chemical processes. As we dive deeper, you'll see how enthalpy relates to bond energies, phase transitions, and even the spontaneity of reactions. So, let's get started with this fundamental concept that's key to unlocking the mysteries of chemical energetics!

Example

Enthalpy changes Exothermic and Endothermic reactions

Step 1: Introduction to Enthalpy and Its Importance in Chemistry

Enthalpy is a fundamental concept in chemistry that deals with the heat content of a system. It is crucial for understanding how energy is transferred during chemical reactions. In this section, we will explore the basic definitions and significance of enthalpy in chemical processes.

Step 2: Understanding Exothermic and Endothermic Reactions

Exothermic and endothermic reactions are two types of chemical reactions that involve heat transfer. Exothermic reactions release heat to the surroundings, resulting in a temperature increase. In contrast, endothermic reactions absorb heat from the surroundings, leading to a temperature decrease. We will delve into the characteristics and examples of each type of reaction.

Step 3: Energy Diagrams and Reaction Coordinates

Energy diagrams are graphical representations that help visualize the energy changes during a chemical reaction. The x-axis represents the reaction coordinate, indicating the progress of the reaction, while the y-axis represents the energy levels of reactants and products. We will examine how these diagrams illustrate the energy changes in exothermic and endothermic reactions.

Step 4: Exothermic Reactions in Detail

In exothermic reactions, the reactants have higher energy than the products. The excess energy is released to the surroundings as heat. This section will provide a detailed explanation of exothermic reactions, including common examples such as combustion and the formation of bonds.

Step 5: Endothermic Reactions in Detail

Endothermic reactions involve reactants with lower energy than the products. These reactions absorb heat from the surroundings to proceed. We will explore the characteristics of endothermic reactions, including examples like photosynthesis and the breaking of bonds.

Step 6: The Role of Bond Enthalpies in Determining Enthalpy Changes

Bond enthalpies play a crucial role in calculating enthalpy changes during chemical reactions. This section will explain how bond enthalpies are used to determine whether a reaction is exothermic or endothermic. We will also discuss the concept of bond breaking and bond formation in relation to energy changes.

Step 7: Calorimetry and Measuring Enthalpy Changes

Calorimetry is a technique used to measure the heat changes during chemical reactions. By monitoring temperature changes in the surroundings, we can determine the enthalpy changes of a reaction. This section will provide an overview of calorimetry and its application in studying exothermic and endothermic reactions.

Step 8: Enthalpy of Formation and Its Significance

The enthalpy of formation is another method to calculate enthalpy changes. It involves determining the heat change when one mole of a compound is formed from its elements in their standard states. We will discuss the importance of enthalpy of formation and how it is used in chemical thermodynamics.

Step 9: Net Heat Changes in Chemical Reactions

Both exothermic and endothermic reactions involve a net change in heat energy. Exothermic reactions result in a net release of heat, while endothermic reactions lead to a net absorption of heat. This section will explain the concept of net heat changes and how it affects the surroundings during chemical reactions.

Step 10: Practical Applications and Examples

Understanding enthalpy changes is essential for various practical applications, including industrial processes, environmental science, and everyday life. We will explore real-world examples of exothermic and endothermic reactions and their significance in different fields.

FAQs

Here are some frequently asked questions about enthalpy:

1. What is the delta H equation?

The delta H equation is ΔH = H(products) - H(reactants). This equation calculates the change in enthalpy for a chemical reaction by subtracting the enthalpy of the reactants from the enthalpy of the products.

2. What does it mean to have a negative enthalpy?

A negative enthalpy (ΔH < 0) indicates an exothermic reaction. This means the system releases heat to its surroundings during the reaction. Exothermic reactions typically feel warm to the touch and can occur spontaneously.

3. What if the enthalpy is positive?

A positive enthalpy (ΔH > 0) signifies an endothermic reaction. In this case, the system absorbs heat from its surroundings. Endothermic reactions often feel cool to the touch and may require energy input to proceed.

4. How do you calculate ΔH?

ΔH can be calculated using various methods, including:

  • Direct measurement using calorimetry
  • Using standard enthalpies of formation
  • Applying Hess's Law for multi-step reactions
  • Using bond enthalpies for gas-phase reactions
The method chosen depends on the available data and the nature of the reaction.

5. Does negative enthalpy mean spontaneous?

Not necessarily. While a negative enthalpy (exothermic reaction) often contributes to spontaneity, it's not the sole determining factor. Spontaneity depends on the Gibbs free energy change (ΔG), which considers both enthalpy and entropy changes. Some endothermic reactions can be spontaneous if there's a significant increase in entropy.

Prerequisites

Understanding enthalpy is a crucial concept in thermodynamics and chemistry, but it's important to recognize that this topic doesn't exist in isolation. While there are no specific prerequisite topics listed for enthalpy in this case, it's worth noting that a strong foundation in basic chemistry and physics principles is essential for grasping the concept of enthalpy fully.

Enthalpy, which is a measure of heat content in a system, builds upon fundamental concepts in thermodynamics. Although we don't have direct links to prerequisite topics, students would benefit from having a solid understanding of energy, heat, and temperature. These basic principles form the groundwork for exploring enthalpy and its applications in various chemical processes.

Additionally, familiarity with the laws of thermodynamics would greatly enhance one's ability to comprehend enthalpy. The first law of thermodynamics, which deals with energy conservation, is particularly relevant when studying enthalpy changes in chemical reactions and physical processes.

Another important aspect to consider is the concept of state functions. Enthalpy is a state function, meaning its value depends only on the current state of the system, not on how it got there. Understanding this property is crucial for grasping how enthalpy behaves in different scenarios.

Moreover, knowledge of chemical bonding and intermolecular forces would provide valuable context for understanding why enthalpy changes occur during chemical reactions or phase transitions. These concepts help explain the energy changes associated with breaking and forming bonds, which are central to enthalpy calculations.

Basic mathematical skills, particularly in algebra and calculus, are also beneficial when studying enthalpy. These mathematical tools are often used in deriving enthalpy equations and solving related problems.

While we don't have specific links to prerequisite topics for enthalpy, it's clear that a strong foundation in general chemistry, physics, and mathematics is crucial. Students who have a good grasp of these fundamental areas will find it easier to understand and apply the concept of enthalpy in various scientific contexts.

In conclusion, although there are no direct prerequisite topics listed here, it's important for students to recognize that enthalpy is part of a broader network of scientific concepts. By strengthening their understanding of related fundamental principles, students can approach the study of enthalpy with greater confidence and insight, leading to a more comprehensive understanding of this important thermodynamic property.