1. The document discusses various types of energy changes that occur during chemical reactions including exothermic reactions, endothermic reactions, and energy level diagrams.
2. It provides examples of heat changes associated with different types of chemical reactions such as heat of precipitation, heat of displacement, heat of neutralization, and heat of combustion.
3. The key concepts covered are how to calculate heat changes using thermochemical equations and energy level diagrams, and how temperature changes can indicate whether a reaction is exothermic or endothermic.
This document discusses various types of thermochemistry and enthalpy changes in chemical reactions. It begins by defining exothermic and endothermic reactions, and describes how enthalpy (ΔH) is the heat released or absorbed during a reaction. It then explains how to calculate enthalpy changes using specific heat and temperature change. The main types of enthalpy covered are heat of precipitation, neutralization, displacement, and combustion. Examples are provided for calculating enthalpy changes using thermochemical equations and experimental temperature changes. In summary, the document provides an overview of thermochemistry concepts and calculations involving enthalpy for different reaction types.
1) Thermochemistry deals with heat changes during chemical reactions. Exothermic reactions release heat while endothermic reactions absorb heat from the surroundings.
2) There are four main types of heat-releasing chemical reactions: precipitation, displacement, neutralization, and combustion. Precipitation involves the formation of an insoluble salt, displacement involves the replacement of one reactant by another, neutralization involves acids and bases reacting to form water, and combustion involves complete burning in oxygen.
3) The heat of reaction, also known as enthalpy change, can be calculated using the formula: Heat of reaction = Q/n where Q is the heat change in Joules and n is the number
This document discusses thermochemistry and energy changes that occur during chemical reactions. It defines exothermic and endothermic reactions, and how to construct energy level diagrams to represent them. Specific heats of reaction like combustion, precipitation, displacement, and neutralization are also explained. Experiments to determine various heats of reaction are described. The relationships between the heat of reaction and type of reactants, as well as the number of carbons in alcohols are also summarized.
This document provides an overview of exothermic and endothermic reactions, calorimetry, and how to calculate enthalpy changes. It discusses key concepts like:
1) The standard enthalpy change (ΔH°) is the enthalpy change measured under standard conditions of temperature and pressure.
2) Calorimetry can be used to determine enthalpy changes by measuring the temperature change of a reaction mixture in a calorimeter.
3) Sample problems demonstrate how to use calorimetry data like temperature changes, masses, and concentrations to calculate enthalpy changes for chemical reactions.
The document discusses thermo chemistry and energy changes in chemical reactions. It covers:
- Exothermic reactions release heat energy to surroundings. Examples include combustion reactions and dissolving acids in water.
- Endothermic reactions absorb heat energy from surroundings. Examples include dissolving salts in water and melting substances.
- The heat of reaction (ΔH) is the energy absorbed or released during a chemical reaction. It is calculated as the energy of products minus reactants.
- Exothermic reactions have a negative ΔH value while endothermic reactions have a positive ΔH.
- Energy diagrams show the relative energy levels of reactants and products.
This document discusses thermo chemistry and key concepts like enthalpy, calorimetry, and Hess's law. It begins by defining open, closed, and isolated systems and explains endothermic and exothermic reactions using energy profile diagrams. It then discusses enthalpy and defines standard enthalpy. Various types of enthalpies are defined like enthalpy of formation, combustion, atomization, and more. Calorimetry is explained as a method to measure heat changes in reactions using calorimeters. Hess's law is introduced as the principle that the enthalpy change of a reaction is the same whether it occurs in one step or multiple steps.
The document discusses heat of displacement reactions. It provides three examples where a more electropositive metal (Zn, Mg) displaces a less electropositive metal (Cu, Pb, Fe) from an aqueous solution. The heat released in each reaction is measured in kJ/mol. It then describes a method to determine the heat of displacement of copper by zinc experimentally. The procedure involves adding zinc to a copper sulfate solution and measuring the temperature change. Based on this experiment, an equation is derived to calculate the heat of displacement of any reaction based on the temperature change and moles of metal displaced.
This document discusses various types of thermochemistry and enthalpy changes in chemical reactions. It begins by defining exothermic and endothermic reactions, and describes how enthalpy (ΔH) is the heat released or absorbed during a reaction. It then explains how to calculate enthalpy changes using specific heat and temperature change. The main types of enthalpy covered are heat of precipitation, neutralization, displacement, and combustion. Examples are provided for calculating enthalpy changes using thermochemical equations and experimental temperature changes. In summary, the document provides an overview of thermochemistry concepts and calculations involving enthalpy for different reaction types.
1) Thermochemistry deals with heat changes during chemical reactions. Exothermic reactions release heat while endothermic reactions absorb heat from the surroundings.
2) There are four main types of heat-releasing chemical reactions: precipitation, displacement, neutralization, and combustion. Precipitation involves the formation of an insoluble salt, displacement involves the replacement of one reactant by another, neutralization involves acids and bases reacting to form water, and combustion involves complete burning in oxygen.
3) The heat of reaction, also known as enthalpy change, can be calculated using the formula: Heat of reaction = Q/n where Q is the heat change in Joules and n is the number
This document discusses thermochemistry and energy changes that occur during chemical reactions. It defines exothermic and endothermic reactions, and how to construct energy level diagrams to represent them. Specific heats of reaction like combustion, precipitation, displacement, and neutralization are also explained. Experiments to determine various heats of reaction are described. The relationships between the heat of reaction and type of reactants, as well as the number of carbons in alcohols are also summarized.
This document provides an overview of exothermic and endothermic reactions, calorimetry, and how to calculate enthalpy changes. It discusses key concepts like:
1) The standard enthalpy change (ΔH°) is the enthalpy change measured under standard conditions of temperature and pressure.
2) Calorimetry can be used to determine enthalpy changes by measuring the temperature change of a reaction mixture in a calorimeter.
3) Sample problems demonstrate how to use calorimetry data like temperature changes, masses, and concentrations to calculate enthalpy changes for chemical reactions.
The document discusses thermo chemistry and energy changes in chemical reactions. It covers:
- Exothermic reactions release heat energy to surroundings. Examples include combustion reactions and dissolving acids in water.
- Endothermic reactions absorb heat energy from surroundings. Examples include dissolving salts in water and melting substances.
- The heat of reaction (ΔH) is the energy absorbed or released during a chemical reaction. It is calculated as the energy of products minus reactants.
- Exothermic reactions have a negative ΔH value while endothermic reactions have a positive ΔH.
- Energy diagrams show the relative energy levels of reactants and products.
This document discusses thermo chemistry and key concepts like enthalpy, calorimetry, and Hess's law. It begins by defining open, closed, and isolated systems and explains endothermic and exothermic reactions using energy profile diagrams. It then discusses enthalpy and defines standard enthalpy. Various types of enthalpies are defined like enthalpy of formation, combustion, atomization, and more. Calorimetry is explained as a method to measure heat changes in reactions using calorimeters. Hess's law is introduced as the principle that the enthalpy change of a reaction is the same whether it occurs in one step or multiple steps.
The document discusses heat of displacement reactions. It provides three examples where a more electropositive metal (Zn, Mg) displaces a less electropositive metal (Cu, Pb, Fe) from an aqueous solution. The heat released in each reaction is measured in kJ/mol. It then describes a method to determine the heat of displacement of copper by zinc experimentally. The procedure involves adding zinc to a copper sulfate solution and measuring the temperature change. Based on this experiment, an equation is derived to calculate the heat of displacement of any reaction based on the temperature change and moles of metal displaced.
The document provides 13 numerical problems related to chemical thermodynamics and energetics. The problems cover topics like work done during gas expansion/compression, standard enthalpy of reaction calculations using bond energies and standard state data, spontaneity of reactions using standard Gibbs free energy and entropy values, heat and enthalpy changes during phase transitions, and calculations involving standard enthalpies of formation.
This document discusses various topics in thermochemistry including:
- Enthalpy changes in chemical reactions and how they are measured using calorimetry. Exothermic and endothermic reactions are explained.
- Hess's law, which states that the enthalpy change of a reaction is independent of the reaction pathway. It can be used to calculate enthalpy changes.
- Standard enthalpies of formation and how they allow calculation of enthalpy changes using Hess's law and bond dissociation enthalpies.
- Measuring enthalpy changes using bomb calorimetry and coffee cup calorimetry. Limitations of each method are discussed.
A reaction can be either fast or slow depending on the time taken. A fast reaction has a short time and high rate, while a slow reaction has a long time and low rate. The rate of reaction depends on factors like concentration, temperature, surface area, and catalysts. Collisions between reactant particles must be effective, with sufficient energy to overcome the activation energy barrier, in order for reactions to occur.
This document discusses energy changes that occur during chemical reactions. It defines endothermic and exothermic reactions, and explains how to interpret energy diagrams for these reactions. Key terms like activation energy, enthalpy change, heat of solution, and heat of neutralization are defined. The document provides examples of how to calculate heat of solution and heat of neutralization using experimental data and equations.
1) Several factors that affect the rate of chemical reactions are discussed, including the size of reactants, concentration of solutions, temperature, and presence of catalysts.
2) When the size of reactants is smaller, the total surface area exposed to collisions increases, leading to a higher rate of reaction.
3) A higher concentration of reactants means more particles per unit volume, resulting in more frequent collisions and a faster rate.
4) At higher temperatures, reactant particles have greater kinetic energy to overcome the activation energy barrier, increasing the frequency of effective collisions and accelerating the rate.
5) Catalysts provide an alternative reaction pathway requiring lower activation energy, allowing more particles to react effectively and speed
The document contains 13 problems related to thermodynamics concepts like heat, work, internal energy, enthalpy, free energy, and chemical equilibrium. The problems involve calculations using these concepts for systems involving gases, reactions, and phase changes. Thermodynamic properties like heat capacities, enthalpies, and free energies are provided to calculate unknown values like temperature changes, energies, and equilibrium constants.
The document defines key thermodynamic terms like heat, work, internal energy, temperature, enthalpy, and heat capacity. It distinguishes between heat and temperature. It discusses specific heat and molar heat capacity, and provides values for common substances. It describes how to calculate heat transfer and enthalpy change using specific heat, molar heat capacity, and temperature change. Hess's law and standard enthalpies of formation are also explained for calculating enthalpy changes in chemical reactions.
The chemical energy of a system is changed as a result of a reaction. Calorimetry. Heat of combustion . Calculation of caloric content of sucrose or food. Combustion reaction.
This document provides an introduction to thermochemistry and the key concepts of enthalpy, enthalpy change, and standard enthalpy of formation. It defines system and surroundings, and the three types of systems - open, closed, and isolated. The key points are:
- Enthalpy change (ΔH) is the difference in enthalpies between products and reactants and indicates whether a reaction is endothermic or exothermic.
- Standard enthalpy of formation (H°f) is the enthalpy change when 1 mole of a substance is formed from its elements under standard conditions.
- Enthalpy of combustion (H°c) is the enthalpy change when 1 mole
2016 topic 5.1 measuring energy changesDavid Young
This document provides an overview of exothermic and endothermic reactions, calorimetry, and how to calculate enthalpy changes. The key points are:
- Exothermic reactions release heat while endothermic reactions absorb heat.
- Enthalpy change (ΔH) is the quantity of heat released or absorbed during a chemical reaction.
- Calorimetry experiments allow calculation of ΔH by measuring the temperature change of a reaction mixture.
- Sample problems demonstrate how to use calorimetry data like mass changes and temperature differences to calculate the enthalpy change for a reaction.
Thermochemistry is the study of heat changes in chemical reactions. There are several types of energy including chemical, thermal, nuclear, and potential energy. Heat is the transfer of thermal energy between objects at different temperatures. Exothermic processes release heat to the surroundings while endothermic processes absorb heat from the surroundings. Enthalpy (H) quantifies the heat flow into or out of a system during chemical reactions at constant pressure. The standard enthalpy of formation (ΔH°f) is the heat change when one mole of a compound forms from its elements. Hess's law states that the enthalpy change is the same whether a reaction occurs in one step or multiple steps.
1. The document contains a practice exam with 37 multiple choice questions covering concepts in thermodynamics and chemistry. The questions cover topics like ideal gases, enthalpy, entropy, spontaneity of reactions, and more.
2. For each question there are 4 possible answers labeled a-d. The correct answers are not provided.
3. The questions are intended to test understanding of fundamental thermodynamic concepts and calculations involving things like heat, work, internal energy, and state functions.
This document discusses fireworks, fire extinguishers, and combustion reactions. It begins with an introduction to how fire and rockets work through various demonstrations. Redox reactions are explained in the context of burning and propulsion. Common gases produced by combustion reactions are identified. The color production mechanisms in fireworks are then covered. Finally, the document discusses different types of fire extinguishers and how to appropriately respond to different types of fires.
The document provides examples of solved problems involving ideal gas laws and gas stoichiometry calculations. The problems cover a range of concepts including determining gas pressures and volumes using the ideal gas equation under different temperature and pressure conditions, calculating mole fractions and partial pressures in gas mixtures, and stoichiometric calculations involving the production and reaction of different gases.
The use of Nauplii and metanauplii artemia in aquaculture (brine shrimp).pptxMAGOTI ERNEST
Although Artemia has been known to man for centuries, its use as a food for the culture of larval organisms apparently began only in the 1930s, when several investigators found that it made an excellent food for newly hatched fish larvae (Litvinenko et al., 2023). As aquaculture developed in the 1960s and ‘70s, the use of Artemia also became more widespread, due both to its convenience and to its nutritional value for larval organisms (Arenas-Pardo et al., 2024). The fact that Artemia dormant cysts can be stored for long periods in cans, and then used as an off-the-shelf food requiring only 24 h of incubation makes them the most convenient, least labor-intensive, live food available for aquaculture (Sorgeloos & Roubach, 2021). The nutritional value of Artemia, especially for marine organisms, is not constant, but varies both geographically and temporally. During the last decade, however, both the causes of Artemia nutritional variability and methods to improve poorquality Artemia have been identified (Loufi et al., 2024).
Brine shrimp (Artemia spp.) are used in marine aquaculture worldwide. Annually, more than 2,000 metric tons of dry cysts are used for cultivation of fish, crustacean, and shellfish larva. Brine shrimp are important to aquaculture because newly hatched brine shrimp nauplii (larvae) provide a food source for many fish fry (Mozanzadeh et al., 2021). Culture and harvesting of brine shrimp eggs represents another aspect of the aquaculture industry. Nauplii and metanauplii of Artemia, commonly known as brine shrimp, play a crucial role in aquaculture due to their nutritional value and suitability as live feed for many aquatic species, particularly in larval stages (Sorgeloos & Roubach, 2021).
The ability to recreate computational results with minimal effort and actionable metrics provides a solid foundation for scientific research and software development. When people can replicate an analysis at the touch of a button using open-source software, open data, and methods to assess and compare proposals, it significantly eases verification of results, engagement with a diverse range of contributors, and progress. However, we have yet to fully achieve this; there are still many sociotechnical frictions.
Inspired by David Donoho's vision, this talk aims to revisit the three crucial pillars of frictionless reproducibility (data sharing, code sharing, and competitive challenges) with the perspective of deep software variability.
Our observation is that multiple layers — hardware, operating systems, third-party libraries, software versions, input data, compile-time options, and parameters — are subject to variability that exacerbates frictions but is also essential for achieving robust, generalizable results and fostering innovation. I will first review the literature, providing evidence of how the complex variability interactions across these layers affect qualitative and quantitative software properties, thereby complicating the reproduction and replication of scientific studies in various fields.
I will then present some software engineering and AI techniques that can support the strategic exploration of variability spaces. These include the use of abstractions and models (e.g., feature models), sampling strategies (e.g., uniform, random), cost-effective measurements (e.g., incremental build of software configurations), and dimensionality reduction methods (e.g., transfer learning, feature selection, software debloating).
I will finally argue that deep variability is both the problem and solution of frictionless reproducibility, calling the software science community to develop new methods and tools to manage variability and foster reproducibility in software systems.
Exposé invité Journées Nationales du GDR GPL 2024
The document provides 13 numerical problems related to chemical thermodynamics and energetics. The problems cover topics like work done during gas expansion/compression, standard enthalpy of reaction calculations using bond energies and standard state data, spontaneity of reactions using standard Gibbs free energy and entropy values, heat and enthalpy changes during phase transitions, and calculations involving standard enthalpies of formation.
This document discusses various topics in thermochemistry including:
- Enthalpy changes in chemical reactions and how they are measured using calorimetry. Exothermic and endothermic reactions are explained.
- Hess's law, which states that the enthalpy change of a reaction is independent of the reaction pathway. It can be used to calculate enthalpy changes.
- Standard enthalpies of formation and how they allow calculation of enthalpy changes using Hess's law and bond dissociation enthalpies.
- Measuring enthalpy changes using bomb calorimetry and coffee cup calorimetry. Limitations of each method are discussed.
A reaction can be either fast or slow depending on the time taken. A fast reaction has a short time and high rate, while a slow reaction has a long time and low rate. The rate of reaction depends on factors like concentration, temperature, surface area, and catalysts. Collisions between reactant particles must be effective, with sufficient energy to overcome the activation energy barrier, in order for reactions to occur.
This document discusses energy changes that occur during chemical reactions. It defines endothermic and exothermic reactions, and explains how to interpret energy diagrams for these reactions. Key terms like activation energy, enthalpy change, heat of solution, and heat of neutralization are defined. The document provides examples of how to calculate heat of solution and heat of neutralization using experimental data and equations.
1) Several factors that affect the rate of chemical reactions are discussed, including the size of reactants, concentration of solutions, temperature, and presence of catalysts.
2) When the size of reactants is smaller, the total surface area exposed to collisions increases, leading to a higher rate of reaction.
3) A higher concentration of reactants means more particles per unit volume, resulting in more frequent collisions and a faster rate.
4) At higher temperatures, reactant particles have greater kinetic energy to overcome the activation energy barrier, increasing the frequency of effective collisions and accelerating the rate.
5) Catalysts provide an alternative reaction pathway requiring lower activation energy, allowing more particles to react effectively and speed
The document contains 13 problems related to thermodynamics concepts like heat, work, internal energy, enthalpy, free energy, and chemical equilibrium. The problems involve calculations using these concepts for systems involving gases, reactions, and phase changes. Thermodynamic properties like heat capacities, enthalpies, and free energies are provided to calculate unknown values like temperature changes, energies, and equilibrium constants.
The document defines key thermodynamic terms like heat, work, internal energy, temperature, enthalpy, and heat capacity. It distinguishes between heat and temperature. It discusses specific heat and molar heat capacity, and provides values for common substances. It describes how to calculate heat transfer and enthalpy change using specific heat, molar heat capacity, and temperature change. Hess's law and standard enthalpies of formation are also explained for calculating enthalpy changes in chemical reactions.
The chemical energy of a system is changed as a result of a reaction. Calorimetry. Heat of combustion . Calculation of caloric content of sucrose or food. Combustion reaction.
This document provides an introduction to thermochemistry and the key concepts of enthalpy, enthalpy change, and standard enthalpy of formation. It defines system and surroundings, and the three types of systems - open, closed, and isolated. The key points are:
- Enthalpy change (ΔH) is the difference in enthalpies between products and reactants and indicates whether a reaction is endothermic or exothermic.
- Standard enthalpy of formation (H°f) is the enthalpy change when 1 mole of a substance is formed from its elements under standard conditions.
- Enthalpy of combustion (H°c) is the enthalpy change when 1 mole
2016 topic 5.1 measuring energy changesDavid Young
This document provides an overview of exothermic and endothermic reactions, calorimetry, and how to calculate enthalpy changes. The key points are:
- Exothermic reactions release heat while endothermic reactions absorb heat.
- Enthalpy change (ΔH) is the quantity of heat released or absorbed during a chemical reaction.
- Calorimetry experiments allow calculation of ΔH by measuring the temperature change of a reaction mixture.
- Sample problems demonstrate how to use calorimetry data like mass changes and temperature differences to calculate the enthalpy change for a reaction.
Thermochemistry is the study of heat changes in chemical reactions. There are several types of energy including chemical, thermal, nuclear, and potential energy. Heat is the transfer of thermal energy between objects at different temperatures. Exothermic processes release heat to the surroundings while endothermic processes absorb heat from the surroundings. Enthalpy (H) quantifies the heat flow into or out of a system during chemical reactions at constant pressure. The standard enthalpy of formation (ΔH°f) is the heat change when one mole of a compound forms from its elements. Hess's law states that the enthalpy change is the same whether a reaction occurs in one step or multiple steps.
1. The document contains a practice exam with 37 multiple choice questions covering concepts in thermodynamics and chemistry. The questions cover topics like ideal gases, enthalpy, entropy, spontaneity of reactions, and more.
2. For each question there are 4 possible answers labeled a-d. The correct answers are not provided.
3. The questions are intended to test understanding of fundamental thermodynamic concepts and calculations involving things like heat, work, internal energy, and state functions.
This document discusses fireworks, fire extinguishers, and combustion reactions. It begins with an introduction to how fire and rockets work through various demonstrations. Redox reactions are explained in the context of burning and propulsion. Common gases produced by combustion reactions are identified. The color production mechanisms in fireworks are then covered. Finally, the document discusses different types of fire extinguishers and how to appropriately respond to different types of fires.
The document provides examples of solved problems involving ideal gas laws and gas stoichiometry calculations. The problems cover a range of concepts including determining gas pressures and volumes using the ideal gas equation under different temperature and pressure conditions, calculating mole fractions and partial pressures in gas mixtures, and stoichiometric calculations involving the production and reaction of different gases.
The use of Nauplii and metanauplii artemia in aquaculture (brine shrimp).pptxMAGOTI ERNEST
Although Artemia has been known to man for centuries, its use as a food for the culture of larval organisms apparently began only in the 1930s, when several investigators found that it made an excellent food for newly hatched fish larvae (Litvinenko et al., 2023). As aquaculture developed in the 1960s and ‘70s, the use of Artemia also became more widespread, due both to its convenience and to its nutritional value for larval organisms (Arenas-Pardo et al., 2024). The fact that Artemia dormant cysts can be stored for long periods in cans, and then used as an off-the-shelf food requiring only 24 h of incubation makes them the most convenient, least labor-intensive, live food available for aquaculture (Sorgeloos & Roubach, 2021). The nutritional value of Artemia, especially for marine organisms, is not constant, but varies both geographically and temporally. During the last decade, however, both the causes of Artemia nutritional variability and methods to improve poorquality Artemia have been identified (Loufi et al., 2024).
Brine shrimp (Artemia spp.) are used in marine aquaculture worldwide. Annually, more than 2,000 metric tons of dry cysts are used for cultivation of fish, crustacean, and shellfish larva. Brine shrimp are important to aquaculture because newly hatched brine shrimp nauplii (larvae) provide a food source for many fish fry (Mozanzadeh et al., 2021). Culture and harvesting of brine shrimp eggs represents another aspect of the aquaculture industry. Nauplii and metanauplii of Artemia, commonly known as brine shrimp, play a crucial role in aquaculture due to their nutritional value and suitability as live feed for many aquatic species, particularly in larval stages (Sorgeloos & Roubach, 2021).
The ability to recreate computational results with minimal effort and actionable metrics provides a solid foundation for scientific research and software development. When people can replicate an analysis at the touch of a button using open-source software, open data, and methods to assess and compare proposals, it significantly eases verification of results, engagement with a diverse range of contributors, and progress. However, we have yet to fully achieve this; there are still many sociotechnical frictions.
Inspired by David Donoho's vision, this talk aims to revisit the three crucial pillars of frictionless reproducibility (data sharing, code sharing, and competitive challenges) with the perspective of deep software variability.
Our observation is that multiple layers — hardware, operating systems, third-party libraries, software versions, input data, compile-time options, and parameters — are subject to variability that exacerbates frictions but is also essential for achieving robust, generalizable results and fostering innovation. I will first review the literature, providing evidence of how the complex variability interactions across these layers affect qualitative and quantitative software properties, thereby complicating the reproduction and replication of scientific studies in various fields.
I will then present some software engineering and AI techniques that can support the strategic exploration of variability spaces. These include the use of abstractions and models (e.g., feature models), sampling strategies (e.g., uniform, random), cost-effective measurements (e.g., incremental build of software configurations), and dimensionality reduction methods (e.g., transfer learning, feature selection, software debloating).
I will finally argue that deep variability is both the problem and solution of frictionless reproducibility, calling the software science community to develop new methods and tools to manage variability and foster reproducibility in software systems.
Exposé invité Journées Nationales du GDR GPL 2024
The debris of the ‘last major merger’ is dynamically youngSérgio Sacani
The Milky Way’s (MW) inner stellar halo contains an [Fe/H]-rich component with highly eccentric orbits, often referred to as the
‘last major merger.’ Hypotheses for the origin of this component include Gaia-Sausage/Enceladus (GSE), where the progenitor
collided with the MW proto-disc 8–11 Gyr ago, and the Virgo Radial Merger (VRM), where the progenitor collided with the
MW disc within the last 3 Gyr. These two scenarios make different predictions about observable structure in local phase space,
because the morphology of debris depends on how long it has had to phase mix. The recently identified phase-space folds in Gaia
DR3 have positive caustic velocities, making them fundamentally different than the phase-mixed chevrons found in simulations
at late times. Roughly 20 per cent of the stars in the prograde local stellar halo are associated with the observed caustics. Based
on a simple phase-mixing model, the observed number of caustics are consistent with a merger that occurred 1–2 Gyr ago.
We also compare the observed phase-space distribution to FIRE-2 Latte simulations of GSE-like mergers, using a quantitative
measurement of phase mixing (2D causticality). The observed local phase-space distribution best matches the simulated data
1–2 Gyr after collision, and certainly not later than 3 Gyr. This is further evidence that the progenitor of the ‘last major merger’
did not collide with the MW proto-disc at early times, as is thought for the GSE, but instead collided with the MW disc within
the last few Gyr, consistent with the body of work surrounding the VRM.
Phenomics assisted breeding in crop improvementIshaGoswami9
As the population is increasing and will reach about 9 billion upto 2050. Also due to climate change, it is difficult to meet the food requirement of such a large population. Facing the challenges presented by resource shortages, climate
change, and increasing global population, crop yield and quality need to be improved in a sustainable way over the coming decades. Genetic improvement by breeding is the best way to increase crop productivity. With the rapid progression of functional
genomics, an increasing number of crop genomes have been sequenced and dozens of genes influencing key agronomic traits have been identified. However, current genome sequence information has not been adequately exploited for understanding
the complex characteristics of multiple gene, owing to a lack of crop phenotypic data. Efficient, automatic, and accurate technologies and platforms that can capture phenotypic data that can
be linked to genomics information for crop improvement at all growth stages have become as important as genotyping. Thus,
high-throughput phenotyping has become the major bottleneck restricting crop breeding. Plant phenomics has been defined as the high-throughput, accurate acquisition and analysis of multi-dimensional phenotypes
during crop growing stages at the organism level, including the cell, tissue, organ, individual plant, plot, and field levels. With the rapid development of novel sensors, imaging technology,
and analysis methods, numerous infrastructure platforms have been developed for phenotyping.
Describing and Interpreting an Immersive Learning Case with the Immersion Cub...Leonel Morgado
Current descriptions of immersive learning cases are often difficult or impossible to compare. This is due to a myriad of different options on what details to include, which aspects are relevant, and on the descriptive approaches employed. Also, these aspects often combine very specific details with more general guidelines or indicate intents and rationales without clarifying their implementation. In this paper we provide a method to describe immersive learning cases that is structured to enable comparisons, yet flexible enough to allow researchers and practitioners to decide which aspects to include. This method leverages a taxonomy that classifies educational aspects at three levels (uses, practices, and strategies) and then utilizes two frameworks, the Immersive Learning Brain and the Immersion Cube, to enable a structured description and interpretation of immersive learning cases. The method is then demonstrated on a published immersive learning case on training for wind turbine maintenance using virtual reality. Applying the method results in a structured artifact, the Immersive Learning Case Sheet, that tags the case with its proximal uses, practices, and strategies, and refines the free text case description to ensure that matching details are included. This contribution is thus a case description method in support of future comparative research of immersive learning cases. We then discuss how the resulting description and interpretation can be leveraged to change immersion learning cases, by enriching them (considering low-effort changes or additions) or innovating (exploring more challenging avenues of transformation). The method holds significant promise to support better-grounded research in immersive learning.
ESPP presentation to EU Waste Water Network, 4th June 2024 “EU policies driving nutrient removal and recycling
and the revised UWWTD (Urban Waste Water Treatment Directive)”
Unlocking the mysteries of reproduction: Exploring fecundity and gonadosomati...AbdullaAlAsif1
The pygmy halfbeak Dermogenys colletei, is known for its viviparous nature, this presents an intriguing case of relatively low fecundity, raising questions about potential compensatory reproductive strategies employed by this species. Our study delves into the examination of fecundity and the Gonadosomatic Index (GSI) in the Pygmy Halfbeak, D. colletei (Meisner, 2001), an intriguing viviparous fish indigenous to Sarawak, Borneo. We hypothesize that the Pygmy halfbeak, D. colletei, may exhibit unique reproductive adaptations to offset its low fecundity, thus enhancing its survival and fitness. To address this, we conducted a comprehensive study utilizing 28 mature female specimens of D. colletei, carefully measuring fecundity and GSI to shed light on the reproductive adaptations of this species. Our findings reveal that D. colletei indeed exhibits low fecundity, with a mean of 16.76 ± 2.01, and a mean GSI of 12.83 ± 1.27, providing crucial insights into the reproductive mechanisms at play in this species. These results underscore the existence of unique reproductive strategies in D. colletei, enabling its adaptation and persistence in Borneo's diverse aquatic ecosystems, and call for further ecological research to elucidate these mechanisms. This study lends to a better understanding of viviparous fish in Borneo and contributes to the broader field of aquatic ecology, enhancing our knowledge of species adaptations to unique ecological challenges.
Authoring a personal GPT for your research and practice: How we created the Q...Leonel Morgado
Thematic analysis in qualitative research is a time-consuming and systematic task, typically done using teams. Team members must ground their activities on common understandings of the major concepts underlying the thematic analysis, and define criteria for its development. However, conceptual misunderstandings, equivocations, and lack of adherence to criteria are challenges to the quality and speed of this process. Given the distributed and uncertain nature of this process, we wondered if the tasks in thematic analysis could be supported by readily available artificial intelligence chatbots. Our early efforts point to potential benefits: not just saving time in the coding process but better adherence to criteria and grounding, by increasing triangulation between humans and artificial intelligence. This tutorial will provide a description and demonstration of the process we followed, as two academic researchers, to develop a custom ChatGPT to assist with qualitative coding in the thematic data analysis process of immersive learning accounts in a survey of the academic literature: QUAL-E Immersive Learning Thematic Analysis Helper. In the hands-on time, participants will try out QUAL-E and develop their ideas for their own qualitative coding ChatGPT. Participants that have the paid ChatGPT Plus subscription can create a draft of their assistants. The organizers will provide course materials and slide deck that participants will be able to utilize to continue development of their custom GPT. The paid subscription to ChatGPT Plus is not required to participate in this workshop, just for trying out personal GPTs during it.
EWOCS-I: The catalog of X-ray sources in Westerlund 1 from the Extended Weste...Sérgio Sacani
Context. With a mass exceeding several 104 M⊙ and a rich and dense population of massive stars, supermassive young star clusters
represent the most massive star-forming environment that is dominated by the feedback from massive stars and gravitational interactions
among stars.
Aims. In this paper we present the Extended Westerlund 1 and 2 Open Clusters Survey (EWOCS) project, which aims to investigate
the influence of the starburst environment on the formation of stars and planets, and on the evolution of both low and high mass stars.
The primary targets of this project are Westerlund 1 and 2, the closest supermassive star clusters to the Sun.
Methods. The project is based primarily on recent observations conducted with the Chandra and JWST observatories. Specifically,
the Chandra survey of Westerlund 1 consists of 36 new ACIS-I observations, nearly co-pointed, for a total exposure time of 1 Msec.
Additionally, we included 8 archival Chandra/ACIS-S observations. This paper presents the resulting catalog of X-ray sources within
and around Westerlund 1. Sources were detected by combining various existing methods, and photon extraction and source validation
were carried out using the ACIS-Extract software.
Results. The EWOCS X-ray catalog comprises 5963 validated sources out of the 9420 initially provided to ACIS-Extract, reaching a
photon flux threshold of approximately 2 × 10−8 photons cm−2
s
−1
. The X-ray sources exhibit a highly concentrated spatial distribution,
with 1075 sources located within the central 1 arcmin. We have successfully detected X-ray emissions from 126 out of the 166 known
massive stars of the cluster, and we have collected over 71 000 photons from the magnetar CXO J164710.20-455217.
3. Exothermic
reaction
- is a chemical reaction that
releases energy to the
surrounding in which the
energy of products is less
than the energy of the
reactants
Endothermic
reaction
- is a chemical reaction that
absorbs energy from the
surrounding in which the
energy of the product is
more than the energy of
reactants
Energy level
diagram
-is a graph that shows energy
change of a chemical reaction
4. Energy changes in chemical reaction
Exothermic Endothermic
Energy is given off
during a chemical
reaction
The surrounding
will become hot
The T of the reaction
mixture will rise
Energy is absorbed
from the surrounding
during a chemical
reaction
The surrounding
will become cold
The T of the reaction
mixture will fall
5. Example of exothermic and
endothermic reaction
Exothermic Endothermic
-photosynthesis
-dissolving ammonium
salts in water
-frying an egg
-decomposition of metals
carbonates
-decomposition of metals
nitrate
-Respiration
-neutralisation
-Reaction acids + metals
-reaction acids + metal
carbonates
-burning of fuels
-dissolving NaOH in water
-Adding H2O to
concentrated acids
-reaction reactive metals
+ H2O
-oxidations of metals
-Haber process
6. Energy level diagrams
Chemical reactions
Certain quantity of heat
is given off or absorbed
(ΔH)
(ΔH) is the difference between
the energy of the reactants and
the energy of the products
(ΔH) = Hproducts – Hreactants
Exothermic reaction
Energy is given off
Hproducts < Hreactants
(ΔH) has a negative sign
Energy
reactants
products
(ΔH) -negative
sign
7. (ΔH) is the difference between
the energy of the reactants and
the energy of the products
(ΔH) = Hproducts – Hreactants
Endothermic reaction
Energy is absorbed
Hproducts > Hreactants
(ΔH) has a positive sign
Energy
reactants
products
(ΔH) positive
sign
8. Interpreting energy level diagrams
Eg.
Energy
Zn + 2HCl
ZnCl2 + H2
ΔH = -126 kJ
Base on the energy level
diagram ,we can say :
The reaction between Zn and HCl to
form ZnCl2 and H2 is exothermic
When one mole of Zn reacts with
two moles of HCl to form one mole
of ZnCl2 and H2, the quantity of heat
released is 126 kJ.
The total energy of one mole of Zn
and two moles of HCl is more than
the total energy of one mole of
ZnCl2 and 1 mole of H2. The
difference in energy is 126 kJ
The temperature of the reaction
mixture will rise
9. Energy
N2 + 2O2
2 NO2
ΔH =+66 kJ
Base on the energy level
diagram ,we can say :
Exercise
The reaction between N2 and O2 to
form NO2 is endothermic
When one mole of N2 gas reacts
with two moles of O2 gas to form
two moles of NO2, the quantity of
heat absorbed is 66 kJ.
The total energy of two moles of
NO2 gas is more than the total
energy of one mole of N2 gas and
two moles of O2. The difference in
energy is 66 kJ
The temperature of the reaction
mixture will fall
10. Energy change and chemical bonds
Refer to your text book –pg 147
Type of
reaction
Exothermic
Endothermic
Energy change Sign of ΔH
Energy absorbed for
bond breaking < energy
released during bond
formation
Energy absorbed for
bond breaking > energy
released during bond
formation
negative
positive
11. Heat of reaction, ΔH is the energy change
of a chemical reaction, that is the
difference between the energy of the
reactants and the energy of the products
The heat of reaction, ΔH is
written at the end of a
chemical equation .
H2 + Cl2 2 HCl ΔH = -184kJ
N2 + O2 2 NO ΔH = +225kJ
These equations called a
thermochemical equation
12. 1. HEAT OF PRECIPITATION
Definition – the energy change when one
mole of precipitate is formed
from its ions
The heat of reaction can be calculated by
using the formula :
Energy change = mcθ
In which, m is the mass of the aqueous
reaction mixture
C is the specific heat capacity of
the aqueous reaction mixture
θ is the change in T
13. Assumptions in this calculation :
1. ρ of the aqueous reaction mixture is
1 g cm-3 ,that is the ρ of water.
2. c of the aqueous reaction mixture is
the same as the specific heat
capacity of water .
The value is 4.2 J g-1 0C-1
3. No heat is lost to or absorbed from
the surroundings
4. No heat is absorbed by the
apparatus of the experiment
14. Determining the heat of
precipitation of AgCl
thermometer
Plastic cup
50 cm3 of
0.5 mol dm-3
AgNO3 solution
50 cm3 of
0.5 mol dm-3
NaCl solution
cover
Reacting
mixture
15. Results
Initial temperature of AgNO3 / 0C
Initial temperature of NaCl / 0C
Average temperature of the
mixture / 0C
Highest temperature of the
mixture/ 0C
Increase in temperature / 0C
28.0
28.0
28.0
32.0
32.0 - 28.0 = 4.0
16. Step 2
Moles of Ag+ = MV/1000
= 0.5x50/1000
= 0.025 mol
Moles of Cl- = MV/1000
= 0.5x50/1000
= 0.025 mol
Calculation :
Step 1
Heat given out
in the reaction = mcθ
= (50 + 50)gx4.2Jg-1 0C-1x40C
= 1680 J
Ag+ + Cl- AgCl
Based on the equation,
1 mole of Ag+will react
with 1 mole of Cl- to
form 1 mole of AgCl.
Therefore , 0.025 mole
of AgCl will formed in
this reaction
17. Step 3
Precipitation of 0.025 mole of AgCl gives out 1680 J
:. Precipitation of 1 mole of AgCl gives out
1680J x 1 mole = 67,200J
0.025 mole
The heat of precipitation of AgCl = -67.2kJ mol-1
(ΔHprecipitation)
Energy
Ag+ + Cl-
AgCl
ΔH = -67.2 kJ mol-1
18. Discussion :
1. The value of heat of precipitation obtained
in the experiment is less than the
theoretical because some heat is lost to
the surroundings .
2. To reduced heat lost to the surroundings:
i) use a plastic cup (good insulator)
ii) mix the AgNO3 solution and NaCl
solution quickly.
19. Exercise 1 – to find the ΔH
100 cm3 of 1 mol dm-3 lead(II) nitrate solution
was mixed with 100 cm3 of 1 mol dm-3
potassium sulphate solution. The temperature
of the reaction rose from 270C to 320C .
Calculate the heat of precipitation of lead(II)
sulphate .(c=4.2 J g-1 0C-1 , density of solution
, 1 gcm-3 )
Solution :
1.Moles of Pb2+ = 1x100/1000 = 0.1 mol
Moles of SO4
2- = 1x100/1000 = 0.1 mol
Pb2+ + SO4
2- PbSO4
- 0.1 mol of Pb2+ react with 0.1 mol of SO4
2-
to form 0.1 mol PbSO4
20. 2. Energy
change = mcθ
= (100 + 100)gx4.2Jg-1 0C-1x(32-27)0C
= 4200J
3.
Precipitation of 0.1 mole of PbSO4 gives out 4200 J
:. Precipitation of 1 mole of PbSO4 gives out
4200J x 1 mole = 42,000J
0.1 mole
The heat of precipitation of PbSO4 = -42kJ mol-1
(ΔHprecipitation)
Pb2+ + SO4
2- PbSO4 ΔH= -42kJ
Exercise – pg 83 (cerdik publication)
21. Exercise 2 – to find θ
The heat of of precipitation of calcium carbonate
Is 12 kJ mol-1. 50 cm3 of 2 mol dm-3 calcium
chloride solution and 50 cm3 of 2 mol dm-3 sodium
carbonate solution are mixed together .Calculate
the temperature change of the reaction mixture .
(c=4.2 J g-1 0C-1 , density of solution, 1 gcm-3 )
Exercise 3-to find the volume of solution
The heat of of precipitation of iron(III) hydroxide
Is -10 kJ mol-1. When a 2 mol dm-3 sodium hydroxide
solution was mixed together with a solution of
iron(III) chloride , 200 J of heat was released .
Calculate the volume of sodium hydroxide solution
that was used .
Ans: 2.860C , 30cm3
22. 2. HEAT OF DISPLACEMENT
Definition – Heat released when one
mole of metal is displaced
from its salt solution by a
more electropositive metal
24. Results
Initial temperature of CuSO4
solution / 0C
Highest temperature of the
mixture/ 0C
Increase in temperature / 0C
28.0
37.0
37.0 - 28.0 = 9.0
25. Step 2
Moles of Cu2+ = MV/1000
= 0.2x50/1000
= 0.01 mol
CuSO4 + Mg MgSO4 + Cu
Based on the equation,
1 mole of CuSO4 will
produced 1 mole of Cu
Calculation :
Step 1
Heat given out
in the reaction = mcθ
= (50)gx4.2Jg-1 0C-1x (37.0-28.0)0C
= 1,890 J
Therefore , 0.01 mole
of Cu2+ will produced
0.01 mole of Cu in
this reaction
26. Step 3
Displacement of 0.01 mole of Cu gives out 1,890 J
:. Displacement of 1 mole of Cu gives out
1,890J x 1 mole = 189,000J
0.01 mole
The heat of diplacement of Cu = -189 kJ mol-1
(ΔHdisplacement)
Energy
Cu2+ +Mg
Mg2+ + Cu
ΔH = -189 kJ mol-1
27. Discussion :
1. The value of heat of displacement obtained
in the experiment is less than the
theoretical because some heat is lost to
the surroundings .
2. To reduced heat lost to the surroundings:
i) use a plastic cup (good insulator)
ii) the metals are added quickly to the
solution.
iii) metals in the powder form are used, so
that the reaction take a shorter time to
complete .
28. Exercise 1 – to find the ΔH
Excess zinc powder is added to 100 cm3 of
1 mol dm-3 iron(III) sulphate solution. The
temperature of the reaction mixture rise by 9.60C .
Calculate the heat of displacement of iron from its
salts solution by zinc . (c=4.2 J g-1 0C-1 , density of
solution, 1 gcm-3 )
Answer = -20.16 kJ mol-1
Exercise – pg 31 (cerdik publication)
29. Relating the volume and the concentration
of solution with temperature change
When excess zinc is added to 50 cm3 of 1.0 mol dm-3
lead(II) nitrate solution,the temperature change
is θ0C. What is the temperature change when
excess zinc is added to :
a)100cm3 of 1.0 mol dm-3 lead(II) nitrate solution ?
b)50 cm3 of 2.0 mol dm-3 lead(II) nitrate solution?
c)100 cm3 of 0.5 mol dm-3 lead(II) nitrate solution?
Try your best !!!
31. Part A
-the no. of moles of Pb(NO3)2 is twice
that of the original solution .
(0.1 mole compared to 0.05 mole)
The heat given off is twice as much
But this heat is distributed over a volume
of solution that is twice as much
(100 cm3 compared to 50 cm3)
The temperature rise is still θ0C
32. Part B
-the no. of moles of Pb(NO3)2 is twice
that of the original solution .
(0.1 mole compared to 0.05 mole)
The heat given off is twice as much
But this heat is distributed over a volume
that is the same
( 50 cm3)
The temperature rise is 2θ 0C
33. Part C
-the no. of moles of Pb(NO3)2 is the
same as that of the original solution .
(0.05 mole)
The heat given off is the same as that
of the original solution.
But this heat is distributed over a volume
that is twice as much
( 100 cm3 compared to 50 cm3)
The temperature rise is θ/2 0C
34. 3. HEAT OF NEUTRALISATION
Definition – the heat
released when one
mole of water is formed
from the neutralisation
between one mole of
hydrogen ions from an acid
and one mole of hydroxide
ions from an alkali .
35. Determining the heat of
neutralisation
thermometer
Plastic cup
50 cm3 of
2.0 mol dm-3
NaOH solution
50 cm3 of
2.0 mol dm-3
HCl acid
cover
Reacting
mixture
36. Results
Initial temperature of NaOH / 0C
Initial temperature of HCl / 0C
Average temperature of acid
and alkali / 0C
Highest temperature of the
mixture/ 0C
Increase in temperature / 0C
28.0
28.0
28.0
40.0
40.0 - 28.0 = 12.0
37. Step 2
Moles of H+ = MV/1000
= 2.0x50/1000
= 0.1 mol
Moles of OH- = MV/1000
= 2.0x50/1000
=0.1 mol
Calculation :
Step 1
Heat given out
in the reaction = mcθ
= (50 + 50)gx4.2Jg-1 0C-1x120C
= 5040J
H+ + OH- H2O
Based on the equation,
1 mole of H+will react
with 1 mole of OH- to
form 1 mole of H2O.
Therefore , 0.1 mole
of H2O will formed in
this reaction
38. Step 3
Formation of 0.1 mole of H2O gives out 5040J
:. Formation of 1 mole of H2O gives out
5040J x 1 mole =50400 J
0.1 mole
The heat of neutralisation = - 50.4kJ mol-1
Between NaOH and HCl
(ΔHneutralisation)
Energy
H+ + OH-
H2O
ΔH = -50.4 kJ mol-1
39. Discussion :
1.Value of heat of neutralisation between a
strong acid and a strong alkali is constant ,
that is -57.3 kJ mol-1
2. The value of heat of neutralisation obtained
in the experiment is less than the
theoretical because some heat is lost to
the surroundings .
3. To reduced heat lost to the surroundings:
i) use a plastic cup (good insulator)
ii) mix the acid and the alkali quickly
40. 4.Value of heat of neutralisation between a
strong acid and a strong alkali is constant ,
that is -57.3 kJ mol-1
.Is it the same if we use :
i) a strong acid and a weak alkali ??
ii) a weak acid and a strong alkali ??
iii) a weak acid and a weak alkali ??
Alert !!!
i) Ethanoic acid is a weak acid and ammonia
solution is a weak alkali ,they both dissociate
partially in water. Most of them still exist as
molecules.
ii) Some of the heat given out during the
neutralisation reaction is used to dissociate
the weak acid or the weak alkali completely in
water . That is why the heat of neutralisation
involving weak acid or weak alkali is less than
-57.3 kJ mol-1
41. Exercise 1 – to find the ΔH
When 100 cm3 of 2 mol dm-3 dilute nitric acid is
added to 100 cm3 of 2 mol dm-3 sodium hydroxide ,
the temperature of the reaction mixture rises from
270C to 40.650C . Calculate the heat of neutralisation
(c=4.2 J g-1 0C-1 , density of solution, 1 gcm-3 )
Answer = -57.33 kJ mol-1
Exercise – pg 219 (cerdik publication)
42. Exercise 2 – to find θ
The thermochemical equation for the reaction
between nitric acid and sodium hydroxide
solution is shown below :
HNO3 + NaOH NaNO3 + H2O ΔH = -57.3 kJ mol-1
When 250 cm3 of 1.0 mol dm-3 nitric acid is added
to 200 cm3 of 2.0 mol dm-3 sodium hydroxide
solution, what is the change in temperature ?
(c = 4.2 J g-1 0C-1 , density of water = 1g cm-3)
Ans: 7.60C
43. Exercise 3 – calculation involving a
dibasic acid
200 cm3 of 2 mol dm-3 sodium hydroxide solution
was added to a fixed volume of 1 mol dm-3 dilute
sulphuric acid . Calculate :
1) the quantity of heat given off
2) the volume of dilute sulphuric acid used .
(c = 4.2 J g-1 0C-1 , density of water = 1g cm-3 ,
ΔH = -57.3 kJ mol-1 )
Ans: 22920J , 200cm3
44. 4. HEAT OF COMBUSTION
Definition – the heat released when one
mole of a substance is burnt
completely in an excess of
oxygen
45. Determining the heat of
combustion
thermometer
water
ethanol
windshield
Copper can
Spirit lamp
Wooden block
46. Results
Initial temperature of 200g
water /0C
Highest temperature of 200
g water / 0C
Increase in the temperature
/ 0C
Mass of lamp and ethanol
before burning / g
Mass of lamp and ethanol
after burning / g
Mass of ethanol burnt / g
28.0
58.0
58.0 – 28.0 = 30.0
245.85
245.85 - 244.95 = 0.90
244.95
47. Step 2
Mass of ethanol
= 245.85 - 244.95
= 0.90g
Moles of ethanol
= 0.9 /molar mass
=0.9g/2(12)+5(1)+16(1)+1(1)gmol-1
=0.0196 mol
Calculation :
Step 1
Heat given out
in the reaction = mcθ
= (200)gx4.2Jg-1 0C-1x300C
= 25200J
48. Step 3
0.0196 mole of C2H5OH gives out 25200 J
:. 1 mole of C2H5OH gives out
25200J x 1 mole =1286000 J
0.0196 mole
The heat of combustion = - 1286 kJ mol-1
of ethanol
(ΔHcombustion)
Energy
C2H5OH + 3O2
2CO2 + 3H2O
ΔH = -1286 kJ mol-1
49. Exercise 1 – to find the ΔH
1.6 g methanol is burnt completely in excess oxygen
.The heat given off is used up to heat up 300 cm3 of
Water . The temperature of the water rises by 28.80C.
Calculate the heat of combustion of methanol
(Ar H = 1 , C=12, O=16 , c=4.2 J g-1 0C-1 , density
of water, 1 gcm-3 )
Answer = -725.76 kJ mol-1
Exercise – pg 141 (cerdik publication)
50. Exercise 2 – to find θ
The thermochemical equation for the combustion
of glucose is shown below :
C6H12O6 + 6 O2 6 CO2 + 6 H2O ΔH = -2400 kJ mol-1
36 g of glucose is burnt completely to heat up
800 cm3 of water. Calculate the rise in
temperature of the water.
(Mr of glucose = 180 , c = 4.2 J g-1 0C-1 , density
of water = 1g cm-3)
Ans: 142.90C
51. Exercise 3 – to find mass of alcohol burnt
Complete combustion of 1 mole of butan-1-ol
produces 2678 kJ of heat . Calculate the mass of
butan-1-ol needed to burn completely in excess
oxygen in order to raise the temperature of
500 cm3 of water by 350C .
(c = 4.2 J g-1 0C-1 , density of water = 1g cm-3)
Ans: 2.03g
52. Relationship between relative molecular
mass and heat of combustion of alcohols
Relative molecular mass
of alcohol increases
Heat of combustion of
alcohol increases
Refer your revision book
53. FUEL VALUE
Definition :
-the amount of heat energy given out when
one gram of the fuel is completely burnt
in excess of oxygen .
-the unit for fuel value if kJ g-1
-a fuel with a high fuel value releases a lot
of heat per gram when it burns.
54. Exercise 4 – to find mass of alcohol burnt
The fuel value of charcoal is 35 kJ g-1 . How much
charcoal must be burnt to boil 1.8 dm3 of water ?
(c = 4.2 J g-1 0C-1 , density of water = 1g cm-3 ,
room temperature of water ,270C )
Ans: 15.77g
55. End of the topic – please come and see
me if you have any problems
Success is simply the natural outcome of
our directed intentions and actions .
SUCCESS DOES NOT COME TO YOU .
YOU GO TO IT.
SUCCESS IS A RESULT , NOT A GOAL
56. Never doubt
You have great
potential within you.
Know yourself and
appreciate all the special
qualities you have.
Never doubt your abilities.
Concentrate on whatever
make you happy and
build your spirit up.
Always aim higher than
you believe you can reach.
Be persistent and consistent.
Do not be afraid to walk
on new path if your
heart lead you to.
Believe in yourself-
that you can ultimately
achieve what you want in life .