7-EnthalpyOfReaction
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This document provides instructions on how to determine the enthalpy of a chemical reaction using calorimetry. Key points include: 1) Calorimetry involves measuring the temperature change of a reaction to calculate the heat released or absorbed. 2) Reactions can be exothermic (releasing heat) or endothermic (absorbing heat). 3) The enthalpy change (ΔH) of a reaction is an intensive property representing the heat of the reaction per mole. 4) Hess's law states that the enthalpy change of a reaction is independent of the pathway and can be determined from the sum of multiple steps.
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Here are the steps to solve this problem using Hess's Law:
1) Write the balanced equation for the reaction of interest:
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- Relationships between internal energy (ΔE), enthalpy (ΔH), heat (q), and work (w)
AP_Chem_Thermodynamics.pptx
Thermodynamics is the study of energy and its transformations. Thermochemistry is the subdiscipline involving chemical reactions and energy changes. The first law of thermodynamics states that energy is conserved as it transforms between forms. Energy exists in kinetic and potential forms. Kinetic energy is associated with motion, while potential energy is stored energy due to position or composition. Heat is the transfer of kinetic energy between objects of different temperatures until they reach equilibrium. Exothermic processes release heat from a system, while endothermic processes absorb heat into a system. Enthalpy changes (ΔH) quantify the heat absorbed or released by chemical reactions. Spontaneous processes are those that proceed without outside intervention, as dictated by a negative change
ch6.pptx
What constitutes waste depends on the eye of the beholder; one person's waste can be a resource for another person.[1] Though waste is a physical object, its generation is a physical and psychological process.[1] The definitions used by various agencies are as below.
United Nations Environment Program
According to the Basel Convention on the Control of Transboundary Movements of Hazardous Wastes and Their Disposal of 1989, Art. 2(1), "'Wastes' are substance or objects, which are disposed of or are intended to be disposed of or are required to be disposed of by the provisions of national law".[2]
United Nations Statistics Division
The UNSD Glossary of Environment Statistics[3] describes waste as "materials that are not prime products (that is, products produced for the market) for which the generator has no further use in terms of his/her own purposes of production, transformation or consumption, and of which he/she wants to dispose. Wastes may be generated during the extraction of raw materials, the processing of raw materials into intermediate and final products, the consumption of final products, and other human activities. Residuals recycled or reused at the place of generation are excluded."
European Union
Under the Waste Framework Directive 2008/98/EC, Art. 3(1), the European Union defines waste as "an object the holder discards, intends to discard or is required to discard."[4] For a more structural description of the Waste Directive, see the European Commission's summary.
Types of Waste
Municipal Waste
The Organization for Economic Co-operation and Development also known as OECD defines municipal solid waste (MSW) as “waste collected and treated by or for municipalities”. [5] Typically this type of waste includes household waste, commercial waste, and demolition or construction waste. In 2018, the Environmental Protection Agency concluded that 292.4 tons of municipal waste was generated which equated to about 4.9 pounds per day per person. Out of the 292.4 tons, approximately 69 million tons were recycled, and 25 million tons were composted. [6]
Household Waste and Commercial Waste
Household waste more commonly known as trash or garbage are items that are typically thrown away daily from ordinary households. Items often included in this category include product packaging, yard waste, clothing, food scraps, appliance, paints, and batteries.[7] Most of the items that are collected by municipalities end up in landfills across the world. In the United States, it is estimated that 11.3 million tons of textile waste is generated. On an individual level, it is estimated that the average American throws away 81.5 pounds of clothes each year.[8] As online shopping becomes more prevalent, items such as cardboard, bubble wrap, shipping envelopes are ending up in landfills across the United States. The EPA has estimated that approximately 10.1 million tons of plastic containers and packaging ended up landfills in 2018. The EPA noted that only 30.
Module 7 - Energy Balance chemical process calculations
Chemical process calculations involve various computations and analyses to design, optimize, and understand chemical processes. Here are some descriptions highlighting different aspects of chemical process calculations:
Material Balances:
Material balances are the cornerstone of chemical process calculations, ensuring that the amounts of all components entering and leaving a system are properly accounted for.
These calculations involve tracking the flow rates and compositions of substances throughout a process, often using mass or mole balances.
Energy Balances:
Energy balances involve quantifying the energy inputs and outputs in a chemical process.
These calculations are crucial for understanding the heat transfer requirements, evaluating energy efficiency, and optimizing process conditions.
Reaction Kinetics:
Chemical reactions kinetics calculations focus on understanding the rates at which reactions occur and how they are influenced by various factors such as temperature, pressure, and catalysts.
These calculations help in determining the optimal reaction conditions and designing reactors for desired conversion rates.
Phase Equilibrium Calculations:
Phase equilibrium calculations deal with determining the distribution of components between different phases in a system, such as liquid-liquid or vapor-liquid equilibrium.
These calculations are essential for designing separation processes like distillation, extraction, and absorption.
Thermodynamic Calculations:
Thermodynamic calculations involve applying thermodynamic principles to predict the behavior of chemical systems.
These calculations include determining properties such as enthalpy, entropy, Gibbs free energy, and fugacity, which are crucial for process design and optimization.
Process Simulation:
Process simulation involves using computer software to model and simulate chemical processes.
These simulations allow engineers to predict process behavior under different operating conditions, optimize process parameters, and troubleshoot potential issues.
7
network error
Ch6 Thermochemistry (updated)
Thermochemistry is the study of heat changes in chemical reactions. There are several types of energy including chemical, thermal, nuclear, and radiant energy. Heat is the transfer of thermal energy between objects at different temperatures. Thermochemistry examines heat absorbed or released by chemical reactions using concepts like exothermic, endothermic, enthalpy, and calorimetry. The first law of thermodynamics states that energy cannot be created or destroyed, only transferred between systems and their surroundings.
Arizona State University 1 School of Molecular Sciences .docx
Arizona State University 1 School of Molecular Sciences
How Can the Thermodynamics of
Dissolution be Defined?
Introduction:
The Gibbs-Helmholtz equation expresses the relationship between the free-energy
change, the enthalpy change, and the entropy change at constant temperature and
pressure:
∆G = ∆H - T∆S Equation 1
From knowing the value of ∆G, you may predict whether a process/reaction will be
spontaneous at a certain temperature. A process is spontaneous if ∆G is negative (∆G <
0), nonspontaneous if ∆G is positive (∆G > 0), and at equilibrium if ∆G = 0.
The enthalpy change, ∆Hrxn, is the heat gained or lost by a system during a reaction
carried out at constant pressure. Most reactions occur in several steps, with energy
required (endothermic, positive ∆H) because energy is needed to break bonds, and
energy released (exothermic, negative ∆H) because energy is released as new bonds
are formed. ∆Hrxn represents the total change in heat energy or enthalpy over the course
of the reaction. In this experiment, you will use a coffee-cup calorimeter to determine the
heat absorbed or released during the dissolution of ammonium chloride and the
dissolution of calcium chloride. From observing the contents of the coffee-cup calorimeter,
you will decide whether the dissolution processes are spontaneous or nonspontaneous.
You will also calculate values of ∆Grxn to check your prediction.
From the law of conservation of energy (energy is conserved) the total energy for the
dissolution process is:
qsystem + qsurroundings = 0 or qsystem = - qsurroudnings Equation 2
where qsystem (or qrxn) represents the heat gained or lost by dissolving the solid (the
system), and qsurroundings (or qsolution) is the heat gained or lost by the solution in the
calorimeter (the surroundings). Thus, heat energy is essentially transferred between the
dissolving solid and the solution in the calorimeter. (For this experiment, the heat
absorbed by the cup, probe, and surroundings can be considered as negligible, so it is
not included in the expression above.)
The heat absorbed or released by the contents of the calorimeter is given by:
q = (mass solution)*(specific heat of solution)*(change in temperature)
q = m· Cs · T, where T=Tfinal-Tinitial Equation 3
The mass of the solution is the sum of the masses of the water and solid placed in the
calorimeter. (Recall that the density of water is1.00 g/mL.) Because the solution is very
Arizona State University 2 School of Molecular Sciences
dilute, the specific heat of the solution is basically equal to that of water, which is 4.184
J/gºC. To calculate change (∆) for a variable it is always final minus initial. The value of
qsurr can be calculated.
Thermochemistry
Thermochemistry is the study of heat changes in chemical reactions. Thermal energy is the energy associated with the random motion of atoms and molecules. Enthalpy (H) is a state function that represents the total energy of a system at constant pressure. The standard enthalpy of formation (ΔHf°) is the change in enthalpy that accompanies the formation of one mole of a compound from its elements in their standard states. Hess's law states that the enthalpy change for a reaction is equal to the sum of the enthalpy changes of the steps in any possible pathway between the initial and final states of the reaction.
Enthalpy
Enthalpy is an important concept in thermochemistry that represents the heat transferred during a chemical reaction under constant pressure. It can be used to determine if a reaction is exothermic or endothermic based on the sign of ΔH. Hess's law states that the enthalpy change of a reaction is independent of the pathway and can be calculated from the enthalpies of the steps. Properties of enthalpy include it being a state function, extensive property, and dependent on the states of the reactants and products.
Energetics1
The document defines standard enthalpy change, exothermic and endothermic reactions, and signs for exo/endothermic reactions. It provides examples of enthalpy calculations, experiments to determine enthalpy changes, and defines combustion enthalpy, formation enthalpy, and Hess's law. Hess's law states that the enthalpy change for a reaction is independent of the pathway and depends only on the initial and final states.
Chemical Reactions.ppt
1) The document discusses two chemical reactions: reaction (1) of CH4 and O2 producing CO2 and H2O, and reaction (2) of CH4 and CO2 producing 2CO and 2H2.
2) Thermodynamic properties such as enthalpy (H) and Gibbs free energy (G) are introduced to explain why reaction (1) is exothermic and proceeds at 400K, while reaction (2) is endothermic and does not proceed at 400K.
3) The concepts of chemical equilibrium are then covered, noting that reactions may stop before reactants are fully consumed, with the ratio of products to reactants (reaction quotient) equaling the equilibrium constant (K
lec 2 dr. marwa.ppsx
This document discusses Gibbs free energy (G) and how it relates to chemical reactions and equilibrium. It defines G and explains how it depends on enthalpy (H), entropy (S), pressure (P), and temperature (T). It describes how the change in G for a reaction (DG) determines whether the reaction is spontaneous. It shows how DG can be used to calculate equilibrium constants (K) at different temperatures and how K values change with temperature.
LabQuest 18 Chemistry with Vernier 18 - 1 Additi.docx
LabQuest
18
Chemistry with Vernier 18 - 1
Additivity of Heats of Reaction:
Hess’s Law
In this experiment, you will use a Styrofoam-cup calorimeter to measure the heat released by
three reactions. One of the reactions is the same as the combination of the other two reactions.
Therefore, according to Hess’s law, the heat of reaction of the one reaction should be equal to the
sum of the heats of reaction for the other two. This concept is sometimes referred to as the
additivity of heats of reaction. The primary objective of this experiment is to confirm this law.
The reactions we will use in this experiment are:
(1) Solid sodium hydroxide dissolves in water to form an aqueous solution of ions.
NaOH(s) Na+(aq) + OH–(aq) H1 = ?
(2) Solid sodium hydroxide reacts with aqueous hydrochloric acid to form water and an aqueous
solution of sodium chloride.
NaOH(s) + H+(aq) ) + Cl–(aq) H2O(l) + Na+(aq) + Cl–(aq) H2 = ?
(3)
Solution
s of aqueous sodium hydroxide and hydrochloric acid react to form water and
aqueous sodium chloride.
Na+(aq) + OH–(aq) + H+(aq) ) + Cl–(aq) H2O(l) + Na+(aq) + Cl–(aq) H3 = ?
In the space below, combine two of the above equations algebraically to obtain the third
equation. Indicate the number of each reaction on the shorter lines.
______ __________________________________________________________________
______ __________________________________________________________________
______ __________________________________________________________________
You will use a Styrofoam cup in a beaker as a calorimeter, as shown in Figure 1. For purposes of
this experiment, you may assume that the heat loss to the calorimeter and the surrounding air is
negligible. Even if heat is lost to either of these, it is a fairly constant factor in each part of the
experiment, and has little effect on the final results. You will make three grand assumptions in
this experiment:
1. The specific heat of solution is equivalent to the specific heat of water.
2. The density of solution is equivalent to the density of water.
3. All solutions are identical in temperature before mixing.
OBJECTIVES
In this experiment, you will
Combine equations for two reactions to obtain the equation for a third reaction.
Use a calorimeter to measure the temperature change in each of three reactions.
Calculate the heat of reaction, H, for the three reactions.
Use the results to confirm Hess’s law.
LabQuest 18
18 - 2 Chemistry with Vernier
MATERIALS
LabQuest 4.00 g of solid NaOH
LabQuest App ring stand
Temperature Probe utility clamp
50 mL of 1.0 M NaOH stirring rod
50 mL of 1.0 M HCl Styrofoam cup
100 mL of 0.50 M HCl 250 mL beaker
100 mL of water
PROCEDURE
Reaction 1
1. Obtain and wear goggles.
2. Connect the Temperature Probe to LabQuest and choose New from the File menu. If you
have an older sensor that does not auto-ID, manually ...
F y b. sc. chemical equilibria
Chemical equilibrium is briefly discussed with following topics:
Free energy change in a chemical reaction. Thermodynamic derivation of the law of chemical equilibrium.
Definition of ΔG and ΔG◦
Le Chatelier’s principle.
Relationships between Kp, Kc and Kx
Chemical thermodynamics
This document provides an overview of chemical thermodynamics, including:
- The first law of thermodynamics which states that change in internal energy equals heat added plus work done.
- The second law of thermodynamics which states that the entropy of the universe increases for spontaneous processes.
- How changes in entropy and free energy determine whether processes are spontaneous, with spontaneous processes favoring higher entropy and more negative free energy.
Notes for Unit 17 of AP Chemistry (Thermodynamics)
Notes for Unit 17 of AP Chemistry (Thermodynamics)
Similar to 7-EnthalpyOfReaction (20)
Module 7 - Energy Balance chemical process calculations
Module 7 - Energy Balance chemical process calculations
Arizona State University 1 School of Molecular Sciences .docx
Arizona State University 1 School of Molecular Sciences .docx
LabQuest 18 Chemistry with Vernier 18 - 1 Additi.docx
LabQuest 18 Chemistry with Vernier 18 - 1 Additi.docx
Notes for Unit 17 of AP Chemistry (Thermodynamics)
Notes for Unit 17 of AP Chemistry (Thermodynamics)
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7-EnthalpyOfReaction
- 1. Determining the Enthalpy of a Chemical Reaction
- 2. Pre-Lab Bring completed pre-lab to class Use net ionic equations Include pre-lab results in introduction section of report Results of pre-lab assignment are your accepted values Use only for determining percent error
- 3. Enthalpy & Heat Enthalpy (H) - Amount of heat change occurring during reaction Exothermic (heat released, DH < 0) Endothermic (heat absorbed, DH > 0) Heat (q) - Amount of energy transferred q=H under constant pressure
- 4. Calorimetry Measurement of heat changes Performed in a calorimeter which prevents heat loss from the system Constant volume calorimeter Constant pressure calorimeter
- 5. Calorimeter Used to prevent the escape of heat One version - Styrofoam cup in a beaker with a cardboard lid
- 6. q and DH q Units of Joules Extensive property Value only true for the experimental amounts of reactants and products DH Units of Joules/mole Intensive property Value based on the moles given by the coefficients of balanced equation
- 7. Relationship of q and DH WATCH the units!!! DH will be expressed as a negative or positive dependant on the direction of heat flow Negative: exothermic, heat released Positive: endothermic, heat consumed
- 8. q = heat, units of Joules m = mass, units of grams s= specific heat of water, units of DT = change in temperature, units of Initial temperature: Right before solutions are mixed Final temperature: Maximum temperature in data q = msDT
- 9. Finding Enthalpy Mix reactants Measure temperature change Calculate heat change Calculate change in enthalpy (DH)
- 10. Experimental Results Enthalpy values for Reactions 1, 2, and 3 Two experimental values for reaction 3 Value from running the reaction Using Hess’ law from reactions 1 & 2 Accepted values are those from pre-lab
- 11. Hess’s Law Enthalpy of a reaction does not depend on number of steps involved If the enthalpy of a set of reactions is known, use combination of reactions to find enthalpy
- 12. Hess’s Law : Example Determine the enthalpy of the reaction shown using the three reactions given. C(gr) + 2S(rh) CS2(l) DH (kJ/mole) C(gr) + O2(g) CO2(g),DH0rxn1= -393.5 S(rh) + O2(g) SO2(g),DH0rxn2= -296.4 CS2(l) + 3O2(g) CO2(g) + 2SO2(g)DH0rxn3= -1073.6 Does this add up to: C(gr) + 2S(rh) CS2(l)?
- 13. C(gr) + 2S(rh) CS2(l) C(gr) + O2(g) CO2(g) S(rh) + O2(g) SO2(g) CS2(l) + 3O2(g) CO2(g) + 2SO2(g) Therefore, rewrite equations C(gr) + O2(g) CO2(g) 2S(rh) + 2O2(g) 2SO2(g) 2SO2(g) + CO2(g) CS2(l) + 3O2(g) C(gr) + 2S(rh) CS2(l) DH0rxn = DH0rxn1 + 2DH0rxn2 - DH0rxn3 = 87.3 kJ/mol 2 x -1 x
- 14. Experiment Tips Rinse and reuse styrofoam cups Save cups for next class Do not poke hole in cups with temperature probe Record exact volumes of each solution
- 15. Experiment Tips Push collect before mixing solutions Values of DT will vary for each of the three parts Press “Autoscale” button in LoggerPro to expand graph Don’t wait until the last minute to complete calculations Allow time to get help, if needed