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Enthalpy

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Intros
Lessons
  1. Enthalpy changes
  2. Exothermic and Endothermic reactions
  3. Mean bond enthalpy and other enthalpy definitions.
  4. Representing enthalpy changes.
  5. Ways to find the enthalpy change.
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Examples
Lessons
  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?
Topic Notes
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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!

Understanding Enthalpy Change (Delta H)

What is Enthalpy Change (Delta H)?

Enthalpy change, commonly denoted as Delta H (ΔH), is a fundamental concept in chemistry that measures the heat energy transferred between a system and its surroundings during a chemical reaction or physical process. It's a crucial parameter in thermodynamics, helping scientists and engineers understand and predict the energy dynamics of various chemical processes.

Defining Delta H in Chemistry

In chemistry, Delta H represents the difference in enthalpy between the products and reactants of a reaction at constant pressure. Mathematically, it's expressed as:

ΔH = H(products) - H(reactants)

This equation, known as the Delta H equation, quantifies the energy change associated with a chemical transformation. It's important to note that enthalpy itself is a state function, meaning its value depends only on the current state of the system, not on how it reached that state.

Calculating Delta H

There are several methods to calculate Delta H:

  • Direct measurement using calorimetry
  • Using standard enthalpies of formation
  • Applying Hess's Law for multi-step reactions
  • Utilizing bond enthalpies for gas-phase reactions

Each method has its applications and limitations, but they all aim to quantify the energy change in a reaction accurately.

Significance of Positive and Negative Enthalpy Values

The sign of Delta H provides crucial information about the nature of a reaction:

  • Negative Delta H (ΔH < 0): Indicates an exothermic reaction
  • Positive Delta H (ΔH > 0): Indicates an endothermic reaction

Understanding whether a reaction is exothermic or endothermic is vital in various fields, from industrial processes to biological systems.

Exothermic Reactions: Releasing Heat

Exothermic reactions release heat to the surroundings, resulting in a negative Delta H. These reactions typically feel warm to the touch and can be spontaneous. Examples include:

  • Combustion of fuels (e.g., burning methane: CH4 + 2O2 CO2 + 2H2O, ΔH = -890 kJ/mol)
  • Neutralization reactions (e.g., HCl + NaOH NaCl + H2O, ΔH -56 kJ/mol)
  • Exothermic phase changes (e.g., water vapor condensing to liquid)

In these reactions, the chemical bonds formed in the products are stronger than those broken in the reactants, resulting in a net release of energy.

Endothermic Reactions: Absorbing Heat

Endothermic reactions absorb heat from the surroundings, leading to a positive Delta H. These reactions typically feel cool to the touch and often require energy input to proceed. Examples include:

  • Photosynthesis (6CO2 + 6H2O C6H12O6 + 6O2, ΔH +2803 kJ/mol)
  • Melting of ice (H2O(s) H2O(l), ΔH = +6.01 kJ/mol)
  • Decomposition of calcium carbonate (CaCO3 CaO + CO2, ΔH = +178 kJ/mol)

In endothermic processes

Interpreting Positive and Negative Enthalpy Values

Understanding enthalpy is crucial in chemistry, as it helps us comprehend the energy changes occurring during chemical reactions. Enthalpy, denoted as H, is a measure of the heat content of a system. When discussing chemical reactions, we often focus on the change in enthalpy, represented as ΔH (delta H). This change can be either positive or negative, each signifying different energy processes within a reaction.

What does positive enthalpy mean?

A positive enthalpy value (ΔH > 0) indicates an endothermic reaction. In these reactions, the system absorbs heat from its surroundings. This means that the products have a higher energy state than the reactants. Examples of endothermic processes include:

  • The melting of ice (H2O(s) H2O(l), ΔH = +6.01 kJ/mol)
  • The decomposition of calcium carbonate (CaCO3(s) CaO(s) + CO2(g), ΔH = +178.3 kJ/mol)

What does negative enthalpy mean?

Conversely, a negative enthalpy value (ΔH < 0) signifies an exothermic reaction. In these cases, the system releases heat to its surroundings. The products have a lower energy state than the reactants. Common exothermic processes include:

  • The combustion of methane (CH4(g) + 2O2(g) CO2(g) + 2H2O(l), ΔH = -890.4 kJ/mol)
  • The formation of water from hydrogen and oxygen (2H2(g) + O2(g) 2H2O(l), ΔH = -571.6 kJ/mol)

Relationship between enthalpy and heat transfer

The enthalpy change directly relates to the heat transferred during a reaction at constant pressure. For endothermic reactions (positive ΔH), heat is absorbed from the surroundings, often resulting in a decrease in temperature of the reaction mixture. In exothermic reactions (negative ΔH), heat is released to the surroundings, typically causing an increase in temperature.

Interpreting enthalpy values in chemical processes

When analyzing chemical reactions, the magnitude of the enthalpy change provides insight into the energy involved. A larger absolute value of ΔH indicates a greater energy change. For instance, comparing the combustion of methane (-890.4 kJ/mol) to the formation of water (-571.6 kJ/mol), we can conclude that the combustion of methane releases more energy.

It's important to note that enthalpy values are typically reported under standard conditions (25°C and 1 atm pressure). These standard enthalpy changes are denoted as ΔH°. When conditions deviate from standard, the actual enthalpy change may differ slightly.

Visual representation: Energy profiles

Energy profiles or reaction coordinate diagrams are useful tools for visualizing enthalpy changes in reactions. These diagrams plot the energy of a system against the progress of a reaction. For an endothermic reaction, the energy profile shows an upward trend, with the products at a higher energy level than the reactants. Exothermic reactions display a downward trend, with products at a lower energy state.

Practical applications of enthalpy values

Understanding enthalpy changes is crucial in various fields:

  • Chemical engineering: Designing efficient chemical processes and heat exchange systems
  • Biochemistry: Studying metabolic pathways and energy transfer in biological systems
  • Materials science: Developing new materials with specific thermal properties
  • Environmental science

    Calculating Enthalpy Change

    Hey there, chemistry enthusiasts! Today, we're diving into the exciting world of enthalpy change calculations. Don't worry if it sounds intimidating I'll break it down for you step by step, and we'll even tackle some practice problems together. Let's get started!

    Understanding Enthalpy Change (ΔH)

    First things first: enthalpy change, often denoted as ΔH (delta H), is the difference in heat content between the products and reactants in a chemical reaction. The formula is simple:

    ΔH = H(products) - H(reactants)

    This is why you'll often hear about "delta h products - reactants" when discussing enthalpy changes. But how do we actually calculate this? Let's explore two main methods.

    Method 1: Using Bond Enthalpies

    One way to calculate enthalpy change is by using bond enthalpies. Here's how it works:

    1. Identify all bonds broken in the reactants (this requires energy, so it's positive)
    2. Identify all bonds formed in the products (this releases energy, so it's negative)
    3. Sum up the energies for breaking bonds
    4. Sum up the energies for forming bonds
    5. Calculate ΔH = (energy to break bonds) + (energy to form bonds)

    Remember, breaking bonds always requires energy (positive), while forming bonds releases energy (negative).

    Method 2: Hess's Law

    Hess's Law states that the total enthalpy change in a reaction is independent of the pathway taken. This means we can break complex reactions into simpler steps and sum their enthalpy changes. Here's how to use it:

    1. Break down the overall reaction into smaller, known reactions
    2. Look up or calculate the ΔH for each step
    3. Sum up all the ΔH values to get the total enthalpy change

    How to Know if Enthalpy is Positive or Negative

    Understanding whether enthalpy delta h is positive or negative is crucial. Here's a quick guide:

    • Positive ΔH: The reaction absorbs heat (endothermic)
    • Negative ΔH: The reaction releases heat (exothermic)

    Step-by-Step Problem Solving

    Let's walk through solving an enthalpy problem:

    1. Read the problem carefully and identify what's given and what's asked
    2. Choose the appropriate method (bond enthalpies or Hess's Law)
    3. Write out the chemical equation if not provided
    4. Apply the chosen method step by step
    5. Calculate the final ΔH value
    6. Don't forget units (usually kJ/mol)

    Practice Problem

    Let's try a problem together! Calculate the enthalpy change for the reaction:

    CH4 + 2O2 CO2 + 2H2O

    Given bond enthalpies (in kJ/mol): C-H = 413, O=O = 495, C=O = 799, O-H = 463

    Solution:

    1. Bonds broken (reactants): 4(C-H) + 2(O=O) = 4(413) + 2(495) = 2652 kJ
    2. Bonds formed (products):

      Applications of Enthalpy in Chemistry

      Enthalpy, a fundamental concept in thermodynamics, plays a crucial role in various fields of chemistry and everyday life. Understanding enthalpy changes is essential for predicting and controlling chemical reactions, optimizing industrial processes, and explaining biological phenomena. This article explores the practical applications of enthalpy in thermochemistry, biochemistry, and industrial processes, highlighting its significance through real-world examples.

      Thermochemistry Applications

      In thermochemistry, enthalpy is a key concept used to study heat changes in chemical reactions. One of the most common applications is in calorimetry, where scientists measure the heat released or absorbed during chemical processes. This information is vital for understanding reaction energetics and predicting spontaneity. For instance, enthalpy measurements help determine the energy content of fuels, allowing engineers to design more efficient combustion engines and power plants.

      Another important application is in the field of materials science. Enthalpy of formation data is used to calculate bond energies and predict the stability of compounds. This knowledge is crucial in developing new materials with specific properties, such as heat-resistant alloys for aerospace applications or energy-storing materials for batteries.

      Biochemistry Applications

      In biochemistry, enthalpy plays a vital role in understanding metabolic processes and enzyme kinetics. The enthalpy changes associated with biochemical reactions help explain how living organisms maintain their energy balance. For example, the enthalpy of ATP hydrolysis is a fundamental concept in bioenergetics, explaining how cells store and utilize energy for various functions.

      Enthalpy is also crucial in studying protein folding and stability. The enthalpy changes during protein denaturation provide insights into the forces that maintain protein structure, which is essential for drug design and understanding diseases related to protein misfolding, such as Alzheimer's and Parkinson's.

      Industrial Process Applications

      In industrial chemistry, enthalpy is a critical factor in process design and optimization. Chemical engineers use enthalpy data to design reactors, heat exchangers, and distillation columns. For instance, in the production of ammonia via the Haber process, understanding the enthalpy changes helps optimize reaction conditions to maximize yield and minimize energy consumption.

      The food industry also relies heavily on enthalpy concepts. In food processing, enthalpy changes are considered when designing cooking, freezing, and preservation methods. For example, the enthalpy of vaporization is crucial in freeze-drying processes used to preserve foods while maintaining their nutritional value and flavor.

      Environmental Applications

      Enthalpy plays a significant role in environmental chemistry and climate science. The enthalpy of vaporization of water, for instance, is crucial in understanding the Earth's water cycle and climate patterns. Climate models use enthalpy data to predict weather patterns and long-term climate changes. Additionally, enthalpy concepts are applied in developing more environmentally friendly refrigerants and assessing the environmental impact of industrial processes.

      Everyday Life Applications

      While often overlooked, enthalpy affects many aspects of our daily lives. In cooking, understanding enthalpy helps explain why certain cooking methods are more effective for specific foods. For example, the high enthalpy of vaporization of water makes steam an efficient cooking medium. In sports, the concept of enthalpy is applied in designing cooling systems for athletes, such as ice vests, which utilize the enthalpy of fusion of ice to absorb heat from the body.

      Pharmaceutical Applications

      In the pharmaceutical industry, enthalpy is crucial in drug development and formulation. The enthalpy of solution helps predict drug solubility and dissolution rates, which are essential for determining bioavailability. Additionally, enthalpy measurements are used in studying drug-target interactions, aiding in the design of more effective medications.

      Energy Storage Applications

      Enthalpy concepts are fundamental in developing energy storage technologies. Phase change materials (PCMs) used in thermal energy storage systems rely on the enthalpy of fusion to store and release heat efficiently. This technology has applications in building temperature control, solar energy storage, and even in smart textiles for temperature regulation in clothing.

      Common Misconceptions about Enthalpy

      Enthalpy is a fundamental concept in thermodynamics, but it often presents challenges for students. Let's address some common misconceptions and errors to help clarify this important topic.

      One frequent misunderstanding is confusing enthalpy with heat. While related, they're not identical. Enthalpy (H) is a state function that represents the total heat content of a system, including both internal energy and the product of pressure and volume. Heat (q), on the other hand, is energy transferred due to temperature differences. Remember, enthalpy change (ΔH) equals heat transferred at constant pressure, but this doesn't mean enthalpy is heat itself.

      Another common error is misinterpreting the sign of ΔH. A negative ΔH doesn't always mean the reaction feels cold to touch. For instance, the combustion of methane has a negative ΔH but produces heat. The negative sign indicates the system releases energy to the surroundings, not that it becomes colder. Conversely, a positive ΔH doesn't necessarily mean the reaction feels hot. It signifies the system absorbs energy from the surroundings.

      Students often struggle with the concept of standard enthalpy of formation (ΔH°f). It's crucial to understand that this refers to the formation of one mole of a compound from its elements in their standard states. A common mistake is forgetting that the ΔH°f of elements in their standard states is zero. For example, the ΔH°f of H2(g) or O2(g) at 25°C and 1 atm is 0 kJ/mol.

      When dealing with Hess's Law, a frequent error is incorrectly manipulating equations. Remember, when you reverse a reaction, you must also change the sign of ΔH. When you multiply a reaction by a factor, you must multiply ΔH by the same factor. For instance, if you need to halve a reaction, you must also halve its ΔH value.

      Many students mistakenly believe that all exothermic reactions are spontaneous. While exothermic reactions (negative ΔH) are often spontaneous, this isn't always the case. Spontaneity depends on the change in Gibbs free energy (ΔG), which considers both enthalpy and entropy changes. Some endothermic reactions can be spontaneous if there's a significant increase in entropy.

      Another misconception is thinking that bond-breaking always requires energy while bond-forming always releases energy. While it's true that breaking bonds requires energy and forming bonds releases energy, the overall enthalpy change of a reaction depends on the balance between these processes. Don't forget to consider both in your calculations.

      Students sometimes struggle with the concept of state functions. Enthalpy is a state function, meaning its change only depends on the initial and final states, not the path taken. This is why we can use Hess's Law to calculate enthalpy changes for reactions that can't be measured directly.

      Lastly, there's often confusion about the relationship between enthalpy and temperature. While increasing temperature generally increases the enthalpy of a substance, the enthalpy change of a reaction isn't always larger at higher temperatures. The temperature dependence of ΔH is determined by the difference in heat capacities of products and reactants.

      Remember, understanding these concepts takes time and practice. Don't be discouraged if you find them challenging at first. With patience and persistence, you'll develop a solid grasp of enthalpy and its applications in chemistry. Keep working through problems, ask questions when you're unsure, and always strive to connect these concepts to real-world phenomena. Your efforts will pay off, leading to a deeper understanding of thermodynamics and its crucial role in chemical processes.

      Conclusion

      In this enthalpy summary, we've explored the fundamental concepts of enthalpy and its significance in thermodynamics. We discussed how enthalpy (H) represents the total heat content of a system and how changes in enthalpy (ΔH) are crucial in understanding chemical reactions. The introduction video provided a visual foundation for grasping these concepts, making it an essential starting point for your enthalpy journey. To solidify your understanding, we encourage you to tackle practice problems and explore additional resources on delta H calculations. Remember, mastering enthalpy is key to excelling in chemistry and related fields. As you continue to engage with this topic, you'll develop a deeper appreciation for its applications in real-world scenarios. With dedication and consistent practice, you'll soon find yourself confidently navigating enthalpy problems and applying these principles to complex understanding chemical reactions. Keep pushing forward, and you'll unlock the power of enthalpy in your scientific endeavors!

    Enthalpy Changes: Exothermic and Endothermic Reactions

    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.

    Prerequisite Topics

    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.

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 A A \, → B \, B but you do know it from A A \, → C \, C and from B B \, → C \, C, you can find A A \, → B \, B out indirectly.
    • Calorimetry (Calorimetry), which measures enthalpy change based on temperature change in a reactions surroundings.