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AQA 1.8 Thermodynamics REVISION

Allery Chemistry・2 minutes read

Chris Harris discusses thermodynamics concepts relevant for A-level chemistry students, highlighting the Born-Haber cycle and key definitions like lattice enthalpy and Gibbs free energy to enhance understanding. He underscores the importance of integrating theoretical knowledge with practical applications while encouraging viewers to utilize available resources for effective revision and practice.

Insights

Chris Harris from A Lurie Chemistry has created a video specifically for A-level chemistry students that covers key thermodynamics concepts in alignment with the AQA specification, providing a structured approach to mastering the subject.
The video includes a wealth of free resources such as revision videos and tutorials on the A Lurie Chemistry YouTube channel, encouraging students to engage with the material and subscribe for ongoing educational content to enhance their learning experience.
The Born-Haber cycle is a central theme in the video, illustrating how to calculate lattice enthalpies using different pathways to form compounds like lithium chloride, emphasizing the importance of understanding energy changes in chemical reactions and the relationship between enthalpy and entropy.
Entropy is discussed as a crucial factor influencing reaction feasibility, with the video explaining that reactions can still occur even if they absorb energy, provided there is a significant increase in disorder, which is quantified using Gibbs free energy to assess whether a reaction is likely to happen.

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Recent questions

What is enthalpy in chemistry?
Enthalpy is a measure of total energy.
How does entropy affect chemical reactions?
Entropy measures disorder, influencing reaction feasibility.
What is the Born-Haber cycle used for?
It calculates lattice enthalpies of ionic compounds.
What defines an exothermic reaction?
An exothermic reaction releases energy, usually heat.
What is Gibbs free energy?
Gibbs free energy determines reaction feasibility and spontaneity.

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Summary

00:00

A Level Chemistry Thermodynamics Explained

Chris Harris from A Lurie Chemistry presents a video focused on the AQA thermodynamics topic, specifically tailored for A-level chemistry students, ensuring all content aligns with the AQA specification.
The video offers a comprehensive range of free resources, including revision videos, whiteboard tutorials, and exam technique videos available on the A Lurie Chemistry YouTube channel, encouraging viewers to subscribe for continued content creation.
PowerPoint slides used in the video can be purchased from Chris's test shop, providing a portable study aid that can be accessed on smartphones or tablets, or printed for use as revision materials.
Key thermodynamic concepts covered include the Born-Haber cycle, Gibbs free energy, and entropy, with a focus on understanding entropy changes and enthalpy in chemical processes.
Definitions crucial for mastering the topic include enthalpy change of formation (one mole of a compound from its elements in standard states), lattice enthalpy of formation (one mole of solid ionic compound from gaseous ions), and lattice enthalpy of dissociation (opposite of formation).
Additional definitions include enthalpy change of dissociation (breaking one mole of bonds in gaseous molecules), first ionization energy (formation of gaseous 1+ ions from gaseous atoms), and enthalpy change of atomization (one mole of gaseous atoms from an element in standard state).
The second ionization energy (formation of gaseous 2+ ions from gaseous 1+ ions) and first electron affinity (formation of gaseous 1- ions from gaseous atoms) are also defined, along with the second electron affinity (formation of gaseous 2- ions from gaseous 1- ions).
The Born-Haber cycle is introduced as a method to calculate lattice enthalpies, which cannot be measured directly, using a visual representation of the process to illustrate different pathways to form lithium chloride from its elements.
The cycle begins with lithium chloride (LiCl) solid and outlines two routes: one directly from elements in standard states and another from gaseous ions, detailing the necessary steps such as atomization of chlorine and ionization of lithium, all of which are endothermic processes requiring energy input.
The video emphasizes the importance of integrating definitions with practical applications in the Born-Haber cycle, encouraging students to practice cycles to reinforce their understanding of thermodynamic concepts.

13:45

Calculating Lattice Enthalpy of Lithium Chloride

Lithium chloride can be formed from gaseous ions, and its lattice enthalpy of formation can be calculated using two methods: directly from elements in their standard states or through a series of steps in a Born-Haber cycle.
The Born-Haber cycle illustrates that the enthalpy change of a reaction remains constant regardless of the path taken, similar to taking different routes to reach the same destination.
In the Born-Haber cycle, energy changes are categorized as endothermic (positive enthalpy change) or exothermic (negative enthalpy change), with upward movements in the cycle representing endothermic processes and downward movements representing exothermic processes.
The lattice enthalpy of formation for lithium chloride is -409 kJ/mol, while the first ionization energy of lithium is 519 kJ/mol, and the first electron affinity of chlorine is -364 kJ/mol.
To calculate the lattice enthalpy, the enthalpy of atomization of chlorine (which is 242 kJ/mol for Cl2) must be halved to account for the formation of half a mole of Cl, resulting in 121 kJ/mol.
The conversion of solid lithium to gaseous lithium requires 161 kJ/mol, and the first ionization energy of lithium is 519 kJ/mol, which is an endothermic process.
The final calculation for the lattice enthalpy of lithium chloride involves summing the relevant energy changes: +364 (electron affinity) - 519 (ionization) - 161 (sublimation) - 121 (atomization) - 409 (formation), resulting in a lattice enthalpy of -846 kJ/mol.
The calculated lattice enthalpy should be negative, confirming that the formation of lithium chloride from gaseous ions is an exothermic reaction, indicating successful bond formation.
The Born-Haber cycle can also be applied to other compounds, such as sodium chloride, where similar steps are followed to calculate missing data points, ensuring to account for the direction of energy changes.
Theoretical and experimental lattice enthalpies may differ based on the ionic character of the compound, with the presence of additional ionization or electron affinity steps affecting the overall energy calculations.

26:46

Ionic Compounds Covalent Character and Dissolution

Ionic compounds are typically classified as either covalent, ionic, or metallic, with ionic compounds formed from metals and nonmetals, but some ionic compounds exhibit covalent character due to distortion in their ionic structure.
Theoretical lattice enthalpies are calculated based on a perfectly ionic model, which assumes spherical ions with evenly distributed charges, but experimental values often differ, indicating the presence of covalent characteristics in the compounds.
The distortion of positive ions can polarize negative ions, leading to covalent character; the greater the polarization, the more covalent character is present in the ionic compound.
A table of ionic compounds shows that experimental lattice enthalpy values are generally higher than theoretical values, indicating covalent character, with lithium chloride being the most purely ionic and magnesium iodide showing significant covalent character due to larger ion sizes and charges.
The enthalpy change of solution refers to the energy change when one mole of an ionic substance dissolves in the minimum amount of solvent, which is measured when no further enthalpy change occurs upon dilution.
To dissolve an ionic compound, the ionic bonds must break (an endothermic process) and new bonds must form between the solvent and the ions (an exothermic process), with the overall enthalpy change determined by the balance of these two processes.
The hydration process involves water molecules surrounding and stabilizing the separated ions, with the strength of the new bonds formed needing to be equal to or greater than the energy required to break the ionic bonds for the substance to dissolve.
The enthalpy change of solution can be calculated using lattice dissociation enthalpy and enthalpy of hydration, employing a cycle similar to Hess's law to account for energy changes during dissolution.
An example calculation for lithium chloride shows that the enthalpy of solution is -37 kJ/mol, indicating an exothermic reaction, as the energy released during hydration exceeds the energy required to break the ionic lattice.
Entropy is introduced as a measure of disorder, with higher entropy indicating greater disorder among particles, and it plays a role in understanding the energy distribution and behavior of substances in solution.

40:43

Understanding Entropy and Reaction Feasibility

The concept of disorder, or entropy, is illustrated through the analogy of a messy bedroom, emphasizing that it is easier to allow disorder to accumulate than to restore order, which requires energy input.
Entropy tends to increase in natural processes, as systems move towards lower energy states, leading to greater disorder over time, similar to how a neglected bedroom becomes messier.
The arrangement of particles in different states of matter is explained: solids have the lowest disorder with particles in fixed positions, liquids have more disorder with particles slightly spaced apart, and gases exhibit the highest disorder with widely spaced particles.
The number of particles also influences entropy; for example, a reaction producing more moles of gas increases entropy due to more possible arrangements of particles, as seen in the reaction of dinitrogen tetroxide (N2O4) producing nitrogen dioxide (NO2), which increases from one mole to two moles.
Reactions can be spontaneous even if they are endothermic (energy-absorbing) if the increase in entropy is significant enough to overcome the unfavorable enthalpy change, demonstrating that entropy can drive reactions despite high energy requirements.
An example of an endothermic reaction is the combination of hydrated barium hydroxide and ammonium chloride, which results in a significant temperature drop of approximately 20 degrees Celsius, demonstrating how entropy can favorably influence a reaction despite it being energetically unfavorable.
The change in entropy (ΔS) can be calculated using the formula ΔS = S(products) - S(reactants), with standard conditions set at 1 mole of substance, 100 kPa pressure, and a temperature of 298 K, measured in joules per Kelvin per mole.
The Gibbs free energy equation (ΔG = ΔH - TΔS) is introduced to determine reaction feasibility, where ΔH is the enthalpy change in joules per mole, T is the temperature in Kelvin, and ΔS is the entropy change in joules per Kelvin per mole; a negative ΔG indicates a feasible reaction.
For the reaction between barium hydroxide and ammonium chloride at 298 K, the calculated ΔG is -12,118 joules per mole, confirming that the reaction is feasible at this temperature due to the negative value.
Temperature can affect reaction feasibility, and various scenarios can be analyzed to see how changes in ΔH and ΔS impact ΔG, reinforcing that for a reaction to be feasible, ΔG must be negative or zero.

54:57

Assessing Chemical Reaction Feasibility Explained

The feasibility of chemical reactions can be assessed using the equation ΔG = ΔH - TΔS, where ΔG represents Gibbs free energy, ΔH is enthalpy, T is temperature in Kelvin, and ΔS is entropy. Understanding this equation is crucial even without specific numerical values.
To analyze reaction feasibility, split the equation into two parts: the enthalpy (ΔH) and the entropy (TΔS). If ΔH is greater than TΔS, the reaction is not feasible; if TΔS is greater, the reaction can be feasible.
Exothermic reactions with positive entropy (ΔH < 0 and ΔS > 0) are always feasible, even at the lowest temperature (0 Kelvin), resulting in a negative ΔG.
Endothermic reactions with negative entropy (ΔH > 0 and ΔS < 0) are never feasible, regardless of temperature, leading to a consistently positive ΔG.
Freezing water is an exothermic process that occurs below 0 degrees Celsius, as it transitions from a liquid to a solid, which is entropically unfavorable.
The decomposition of sodium hydrogen carbonate (baking soda) is an endothermic process that becomes feasible above a certain temperature. For this reaction, the minimum temperature required for feasibility can be calculated using T = ΔH / ΔS, resulting in a temperature of 385 Kelvin.
The video provides additional resources, including AQA chemistry videos and slides for purchase, encouraging viewers to subscribe for ongoing educational content.

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