Off-Label Drugs in Pediatrics Discussion

Off-Label Drugs in Pediatrics Discussion
This is just a discussion. I will upload an example to use as a guide. Only 550 words needed. Added extra page due to discussion sometimes go over 550 words. Children deal with variety of health issues similar to adults, but they also have issues that are more prevalent within their population. One issue that significantly impacts children is the prescription of drugs for off-label use. Consider the case of Rebecca Riley. When she was two years old, Riley was diagnosed with ADHD, and by age three, she was diagnosed with bipolar disorder. In the span of two years, Riley’s doctor prescribed four drugs off-label: Clonidine, Depakote, Zyprexa, and Seroquel. Riley’s doctor also approved 13 increases in drug dosages. Then, at age four, Riley died from pneumonia combined with a toxic level of prescription drugs (Lambert, 2010). Cases such as this have brought attention to the off-label use of drugs in pediatric patients, as well as the importance of monitoring patient reactions to prescribed drugs and evaluating the effects of drug-drug interactions. As an advanced practice nurse, how do you determine the appropriate use of off-label drugs in pediatrics? Are there certain drugs that should be avoided with pediatric patients? This week you examine the practice of prescribing off-label drugs to children. You also explore strategies for making off-label drug use safer for children from infancy to adolescence as it is essential that you are prepared to make drug–related decisions for pediatric patients in clinical settings.  Off-Label Drugs in Pediatrics Discussion Discussion: Off-Label Drug Use in Pediatrics

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The unapproved use of approved drugs, also called off-label use, with children is quite common. This is because pediatric dosage guidelines are typically unavailable since very few drugs have been specifically researched and tested with children. When treating children, prescribers often adjust dosages approved for adults to accommodate a child’s weight. However, children are not just “smaller” adults. Adults and children process and respond to drugs differently in their absorption, distribution, metabolism, and excretion. Children even respond differently during stages from infancy to adolescence. This poses potential safety concerns when prescribing drugs to pediatric patients. As an advanced practice nurse, you have to be aware of safety implications of the off-label use of drugs with this patient group. To prepare: • Review the Panther et al (2017) and Corney, Lebel, Bailey, and Bussieres (2015) articles in the Learning Resources. Reflect on situations in which children should be prescribed drugs for off-label use. • Think about strategies to make the off-label use and dosage of drugs safer for children from infancy to adolescence. Consider specific off-label drugs that you think require extra care and attention when used in pediatrics. With these thoughts in mind: By Day 3 Post an explanation of circumstances under which children should be prescribed drugs for off-label use. Then, describe strategies to make the off-label use and dosage of drugs safer for children from infancy to adolescence. Include descriptions and names of off-label drugs that require extra care and attention when used in pediatrics. Off-Label Drugs in Pediatrics Discussion

 
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Assignment 1: The Annotated Bibliography

Assignment 1: The Annotated Bibliography

Instructions

Assignment 1: The Annotated Bibliography
Objective: Assess sources for your research for your final presentation (for credibility, reliability, and relevance) and list references in proper APA format
Assignment Instructions: The Research Project/Presentation for this class is divided into three major Assignments, 1) annotated bibliography, 2) outline and 3) final presentation. The first part is the annotated bibliography. An annotation is a summary and evaluation, and your annotated bibliography will include a summary and evaluation of some of the sources (or references) you will use for your presentation.Assignment 1: The Annotated Bibliography

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To prepare for this assignment, I recommend that you do the following:

  • Read these directions carefully.
  • Review the sample annotated bibliography provided to you below.
  • Message me using Classroom Support with any questions!

The reason the annotated bibliography is included as part of the research project is that writing an annotated bibliography is important in that it provides excellent preparation for the final presentation. One of the issues regarding any type of research, especially in chemistry, is the credibility of the sources used, particularly those obtained from various websites. By forcing you to evaluate each of your potential sources carefully, the annotated bibliography helps you determine if in fact the source you chose is credible and helps you determine how relevant it is to your topic and understand the topic better which will help you develop your presentation.
For this project, you will assess three sources to include:
1) a complete citation for each source,
2) a summary of each source, and
3) an evaluation of each source.
Three sources are required for this assignment (i.e., you are to write an annotation for each source). However, you must use five or more sources in your final presentation.Assignment 1: The Annotated Bibliography
Use this TEMPLATE to summarize and evaluate each of your three sources.

  1. Citation:

Written in APA reference list format. For more help with formatting, see APA handout.

  1. Summary:

What is the purpose of the source, review article, original research? What topics are covered? This section is generally 4-6 sentences that summarize the author’s main point. For more help, see this link on paraphrasing sources.

  1. Evaluation:

After summarizing the article (or research paper or book), it is necessary to evaluate it and state where you found it – its source (e.g., journal, website, etc.). Briefly answer the following questions in 4-6 sentences:
What is the format or type of source (e.g., peer-reviewed journal paper, website, book)? How reliable is the information in the article, and how credible is the source (e.g., website’s sponsoring organization, journal or book publisher) and the author(s)?
For more help, see this handout on evaluating resources.
Additional Resources:

  • Sample Annotated Bibliography
  • Also, please see the resources below at The Owl at Purdue site for more information on how to write an annotated bibliography as well as other pages on the site to assist you with the other parts of the research paper:

Assignment 1: The Annotated Bibliography

 
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Hess Law Report Lab

Hess’ Law

Peter Jeschofnig, Ph.D.

Version 42-0158-00-01

Review the safety materials and wear goggles when

working with chemicals. Read the entire exercise

before you begin. Take time to organize the materials

you will need and set aside a safe work space in

which to complete the exercise.

Experiment Summary:

Students will have the opportunity to measure

temperature changes taking place in a calorimeter

during neutralization reactions and use the

measurements to calculate enthalpy of reaction.

They will illustrate the validity of Hazy’ Law by

comparing the values of enthalpy of two chemical

reactions.

Objectives

To measure temperature changes taking place in a calorimeter during neutralization reactions

and use the measurements to calculate enthalpy of reaction.

To compare the enthalpy of two chemical reactions and use these measured values to illustrate

the validity of Hess’ Law.

Materials

Materials From: Label or

Box/Bag: Qty Item Description:

Student Provides Distilled water

Watch

Coffee cups

Paper towels

From LabPaq 1 Thermometer – Digital

1 Goggles-Safety

4 Cup, Styrofoam, 8 oz

1 Cylinder-25-mL

From Experiment Bag

Hess’ Law 2 Ammonia , NH3 (comes as aqueous

ammonia, NH4OH), – 2 M – 10 mL

2 Ammonium chloride, NH4Cl – 2M – 10mL

2 Hydrochloric acid, HCl – 2 M – 20 mL

2 Pipet, Long Thin Stem

2 Sodium hydroxide, NaOH – 2M – 20 mL

Note: The packaging and/or materials in this LabPaq may differ slightly from that which is listed

above. For an exact listing of materials, refer to the Contents List form included in the LabPaq.

Discussion and Review

Thermochemistry is the study of the heat energy involved in chemical reactions and changes of physical state. Nearly all chemical reactions involve the release or absorption of heat, a form of energy. The burning of any fuel such as gasoline, coal, or wood is an example of a heat-releasing reaction. Heat energy is called thermal energy, and it is always spontaneously transferred from hotter to colder matter.

The First Law of Thermodynamics is the Law of Energy Conservation. It states that the total energy of the universe must remain constant. Therefore, all energy transferred between a system and its surroundings must be accounted for as heat or work.

The standard S.I. unit for heat energy is the joule, J. It takes 4.184 joules, the equivalent of 1

calorie, to raise the temperature of one gram of water by 1° C. The kilojoule, kJ, is commonly used in many applications: 1000 joule = 1 kilojoule.

When a chemical reaction takes place in a stable environment where the temperature and

pressure remain constant, the system defined by the reactants and products either produces or

releases heat energy.

If the reacting system releases heat energy to its surroundings, a concurrent increase in

surroundings temperature is observed, and the reaction is exothermic

If the system absorbs heat energy from its surroundings, a decrease in the surroundings

temperature is observed, and the reaction is endothermic.

A measure of the amount of heat given off or absorbed in any chemical reaction is called the

enthalpy change or heat of reaction, and is given the symbol H.

When thermodynamic measurements are carried out at standard-state conditions where the

pressure is constant at 1 atm and the temperature is constant at 25oC, the reaction enthalpy is

designated as the standard enthalpy change or Δ. It is important to have standardized values because the enthalpy of a reaction can vary with different reaction conditions.

The following reaction for the formation of water from its constituents is exothermic:

H2(g) + ½ O2(g) à H2O(l); ΔH °f = -286 kJ

For every mole of H2O (l) formed at standard-state conditions, 286 kilojoules of heat energy are

released. When the standard enthalpy change of reaction describes the formation of 1 mol of

compound directly from its elements in their standard states as in this example, the value of ΔH of is called the standard heat of formation.

To determine the enthalpy change for a given reaction (ΔH°rxn), the summation of the heats of

formation (ΔH° f ) for the reactants are subtracted from the summation of the heats of formation ( ΔH ° f ) for the products.

ΔH° rxn = [n ΔH°f (products)] – [n ΔH°f (reactants)]

Tables containing the standard heats of formation for a number of compounds are available in the appendices of any general chemistry textbook.

Hess’s Law states that if a reaction is the sum of two or more other reactions, the ΔH for the

overall process must be the sum of the ΔH values of the constituent reactions.

Enthalpy change (ΔH) is independent of the path that a reaction follows to move from reactants

to products. It only depends on the relative energy difference between the reactant and product

molecules at constant pressure. Enthalpy change is referred to as a state function due to its

independent of pathway. Since the enthalpy of a substance is not commonly determined, the

change in enthalpy when reactants are converted to products is often used to describe a chemical

or physical process.

The thermal energy absorbed or produced by a chemical process reflects a difference between

the enthalpy between the reactants and products (ΔH). For example, in the decomposition of

liquid water into its component elements, H2 (g) and O2 (g), there are two successive changes.

First, the liquid water is vaporized. Second, the water vapor decomposes into its constituent

elements shown below. The ΔH value for this overall process can be determined by adding the

ΔH values from the equations for each step as shown below.

(1) H2O (l) àH2O (g); ΔH 1 = +44 kJ

(2) H2O (g) àH2 (g) + ½ O2 (g); ΔH 2 = +242 kJ

_______________________________________________________________

(1) + (2) H2O (l) àH2 (g) + ½ O2 (g); ΔHnet = +286 kJ

In order to determine ΔH for the reaction NH3 + HCl àNH4Cl in this experiment, ΔH rxn for the

following two reactions will be measured:

1. NaOH (aq) + HCl (aq) àH2O (l) + NaCl (aq)

2. NaOH (aq) + NH4Cl (aq) àNH3 + NaCl + H2O (l)

Comparison of the calculated results for different parts of the experiment will verify the

generalization known as Hess’s Law of Constant Heat Summation. In this case the target reaction NH3 + HCl àNH4Cl can also be performed directly and the results compared to reactions 1 and 2.

A Styrofoam coffee cup calorimeter will be used to measure the amount of heat energy evolved

or absorbed during the chemical reactions of this experiment. A digital thermometer is used to

measure the change in temperature between the final and initial temperatures of the solutions.

Unfortunately, it is impossible to have perfect insulation and some of the heat energy will be lost to the surroundings, including to the material from which the calorimeter is constructed.

Calibrating the calorimeter before using it to make measurements on an unknown system usually solves the problem of heat losses. A known amount of heat energy from a known process is released into the calorimeter system, and the temperature change is measured. A simple calculation is done to determine the amount of heat energy loss, called the heat capacity of the calorimeter or calorimeter constant. For this experiment it assumed that the heat capacity of the calorimeter is insignificant and it is ignored.

Another practical problem is that heat energy exchanges do not occur instantaneously; i.e., it takes time for energy to move from a hot object to a cold one. An acceptable solution to this problem is to obtain a cooling curve for the heat energy exchange in question and then extrapolate the data back to the exact time that the exchange began.

Below is a sample graph from hypothetical data. Notice that at the time of combining the

two solutions, their starting temperature is 20oC. Since the starting temperatures are at room

temperature no initial temperature adjustment is needed. From 0 to 40 seconds the temperature

rises rapidly to 34.2oC. The temperature then drops gradually 31.1oC and will continue to drop.

Usually recording the temperature in 20-20 second intervals for 5 minutes is enough to provide a

good cooling curve. Extrapolation of these data backward in time determines what the temperature

at the time of mixing would have been if the temperature of the reaction had been instantaneous

and the calorimeter had warmed instantaneously. In this example, the temperature at the time

of mixing determined by extrapolation is 34.3oC.

Calculations: The equation used to calculate heat gained or lost is:

qsolution = (mass of solution) x (specific heat) x ΔT

Density = 1.02 g/mL for all solutions in this experiment;

Specific Heat = 4.184 J

ΔT = Final temperature – Initial temperature

A small amount of heat is lost to the surroundings which in this case is the calorimeter. This

heat loss can be accounted for by using a calorimeter constant, c, which can be determined

experimentally. However, the amount of heat lost to the calorimeter is so insignificant that it is

often left off, or simply assumed to be 1 J* ΔT. (q cal = c x ΔT).

If a correction was to be made for the heat absorbed by the calorimeter, the heat of the reaction,

qrxn , could be determined by taking the negative of the heat gained by the solution, qsoln, plus that

gained by the calorimeter, qcal:

qrxn = -(qsoln + qcal)

Once the total thermal energy transfer is known, the enthalpy of reaction can be determined

using the following equation:

ΔH = qrxn /moles NaOH or HCl

Moles of NaOH or HCl can be determined from the equation: M = moles/L

10 mL = 0.01L; 2M = moles/0.01L = 0.02 moles

Exercise 1: Hess’ Law

Procedure

Part 1: Reaction: HCl & NaOH → NaCl + H2O

1. Before beginning, set up data tables similar to the Data Tables 1 & 2 in the Lab Report Assistant

section.

2. Construct a calorimeter from 2 Styrofoam cups: Trim the lip of one cup and use that cup as

the top of the calorimeter. Make a small hole in the top so a thermometer can be inserted, as

shown below. Be careful when inserting the thermometer into the calorimeter since it has a

pointed tip that could puncture the lower cup if inserted too forcefully. Place the calorimeter

assembly into an empty coffee cup to help prevent it from tipping over.

Figure 2:

3. Use a graduated cylinder to accurately measure 10 mL of 2M HCl. Use an empty thin-stem

pipet to remove or add drops of HCl so that the meniscus level is on the 10 mL mark. Pour the

10 mL HCl into the Styrofoam calorimeter. Rinse the thin-step pipet according to this manual’s

instructions on Use, Disposal, and Cleaning of Common Materials.

4. Rinse and dry the graduated cylinder and accurately measure 10 mL of 2M NaOH using the

same technique in step 2 above. Pour the 10 mL NaOH into another Styrofoam cup and place

the cup into a second empty coffee cup to prevent it from tipping over.

5. Turn on the digital thermometer and place it into the HCl solution. Wait 5 minutes and record

the temperature of the solution in Data Table 1.

6. Remove the thermometer, rinse the tip with distilled water, dry it with a paper towel and

place it into the NaOH solution. Wait 5 minutes and record the temperature of the solution

in Data Table 1. Remove the thermometer, rinse the tip with distilled water, and dry it with a

paper towel for future use.

7. Pour the contents of one Styrofoam cup into the second one, combining the two solutions.

Quickly place the Styrofoam lid on top of the cup containing the combined solutions and insert

the thermometer through the hole in the lid. Be careful when inserting the thermometer to

ensure its pointed tip does not puncture the lower Styrofoam cup.

8. Record the temperature every 20 seconds for 5 – 6 minutes and record in Data Table.

9. Graph the data points using an Excel spreadsheet; time in seconds on the x-axis and

temperature on the y-axis. The graph should look similar to the sample cooling curve below.

10. Place a ruler on the declining temperature portion of the curve and extrapolate to the 0-line.

Read the extrapolated temperature where the straight line intersects the 0-time line. This

temperature represents the final temperature of the mixture. Enter this temperature in Data

Table 1.

11. Dispose of the solution in the calorimeter by flushing it down the drain with water. Recall that

the solution results from a neutralization reaction and is simply salt water.

12. Rinse all equipment used in preparation for reaction 2. This includes the calorimeters,

graduated cylinders, pipets, etc.

Part 2: Reaction 2: NH4Cl + NaOH → NH3 + NaCl + H2O

1. Repeat the Procedures from Part 1, but using 10 mL of 2M NH4Cl and 10 mL of 2 mL of NaOH.

2. Dispose of the solution in the calorimeter by flushing it down the drain with water.

3. Rinse all equipment used in preparation for reaction 3. This includes the calorimeters,

graduated cylinders, pipets, etc

Part 3: Reaction: NH3 + HCl → NH4Cl

1. Repeat the Procedures from reaction 1, but using 10 mL of 2M NH3 and 10 mL of 2 mL of HCl.

2. Dispose of the solution in the calorimeter by flushing it down the drain with water.

3. Rinse all equipment used in preparation for future experiments. This includes the calorimeters,

graduated cylinders, pipets, etc.

Hess’ Law

Peter Jeschofnig, Ph.D.

Version 42-0158-00-01

Lab Report Assistant

This document is not meant to be a substitute for a formal laboratory report. The Lab Report

Assistant is simply a summary of the experiment’s questions, diagrams if needed, and data tables

that should be addressed in a formal lab report. The intent is to facilitate students’ writing of lab

reports by providing this information in an editable file which can be sent to an instructor.

Part 1: Reaction: HCl & NaOH → NaCl + H2O

Part 1: Reaction: HCI & NaOH  →  NaCI +H20

Data Table 1: Sample Data
InitialTemperature of HCl –oC 
InitialTemperature NaOH – oC 
Average InitialTemperature – oC 
Final Temperature of mixture (extrapolated) 
Change in Temperature of mixture, ΔT 
Data Table 2: Sample Data
Time after mixing- seconds Temperature – °C
20 
40 
60 
80 
100 
120 
140 
160 
180 
200 
220 
240 
260 
280 
300 

Part 2: Reaction 2: NH4Cl + NaOH → NH3 + NaCl + H2O

Data Table 3:
InitialTemperature of NaOH – oC 
InitialTemperature NHCl – oC 
Average InitialTemperature – oC 
Final Temperature of mixture (extrapolated) 
Change in Temperature of mixture, ΔT 
Data Table 4:
Time after mixing- seconds Temperature – °C
20 
40 
60 
80 
100 
120 
140 
160 
180 
200 
220 
240 
260 
280 
300 

Part 3: Reaction : NH3 + HCl → NH4Cl

Data Table 5:
InitialTemperature of HCl – oC 
InitialTemperature NH  – oC 
Average InitialTemperature –oC 
Final Temperature of mixture (extrapolated) 
Change in Temperature of mixture, ΔT 
Data Table 6:
Time after mixing- seconds Temperature – °C
20 
40 
60 
80 
100 
120 
140 
160 
180 
200 
220 
240 
260 
280 
300 

Questions

For A. through E. See the calculations for the Data Tables above.

A.      Using the data from your data tables calculate ΔT for all three reactions:

B.      Calculate the heat loss or gain of the three solution mixtures:

C. Use Hess’ Law and ΔH for the first two reactions:

NaOH (aq) + HCl (aq) → H2O (l) + NaCl (aq)

NaOH (aq) + NH4Cl (aq) → NH3 + NaCl + H2O (l)

to determine ΔH for this reaction: NH3 + HCl → NH4Cl

D. Compare the results of step 3 above with the experimental results of the

NH3 + HCl → NH4Cl

E. Use the thermodynamic quantities given below to calculate the theoretical ΔH for this

reaction: NH3 + HCl → NH4Cl

ΔH°f for NH3 (aq) = – 80.29 kJ/mol

ΔH°f for HCl (aq) = – 167.2 kJ/mol

ΔH°f for NH4 (aq) = – 132.5 kJ/mol

ΔH°f for Cl- (aq) = – 167.2 kJ/mol

F. What was the ΔH value obtained for NH3 + HCl àNH4Cl from Hess’ Law method?

G. What was the ΔH value obtained for NH3 + HCl àNH4Cl experimentally?

H. What was the calculated ΔH value obtained for NH3 + HCl àNH4Cl using published

thermodynamic data?

What was the % error of the various methods used? (i.e. comparing the results of the results of Hess’ Law method and the experimental results to the calculated value?

J. Name three examples of the practical application for the use of ΔH values.

 
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Net Ionic Eqaution

Results

Part I: Changes in Reactant or Product Concentrations
A. Copper and Nickel Ions
Observed colors of the solutions:
Copper (II) IonsNickel (II) Ions
CuSO4(aq):clear, light blueNiCl2(aq):clear, light teal
[Cu(NH3)4]2+(aq):clear, dark blue[Ni(NH3)6]2+(aq):cloudy (precip), dark blue/green
After HCl addition:produced heat and formed water vaporAfter HCl addition:produced heat, water vapor, cleared the color, and removed the precipitate
a.) Explain your observations upon addition of NH3(aq) to the CuSO4 and NiCl2 solutions. Include balanced net ionic equations (2) and explain what happens in the reactions in terms of LeChâtelier’s principle as more NH3(aq) is added.b.) Explain your observations upon addition of HCl(aq) on the solutions. Include the balanced net ionic equation for the reaction that occurs. Explain which way the equilibria in (a) shift and why in terms of LeChatlier’s principle as HCl is added.c.) What initially forms as pale blue and pale green precipitates when NH3(aq) is added to [Cu(H2O)4]2+ and [Ni(H2O)6]2+ solutions, respectively? Why do these precipitates form?
B. Cobalt Ions
Color of original CoCl2 solution:Clear, see through red, no precipitate
Color after adding HCl(aq):Dark purple, clear, formed water vapor
Color after adding H2O:Cleared changes, returned to red, see through.
a.) Explain your observation upon addition of 12 M HCl(aq) to the original CoCl2 solution. Include a balanced net ionic equation and explain which way the equilibrium shifts as more HCl(aq) is added and why in terms of LeChâtelier’s principle.b.) Account for your observation when adding water to the [CoCl4]2- complex solution. Explain which way the equilibrium in (a) shifts as water is added and why in terms of LeChâtelier’s principle.
Part II: Equilibria Involving Sparingly Soluble Salts
Appearance of Na2CO3 Solution:Clear, no color
Appearance of AgNO3 Solution:Clear no color
Observed Changes on Mixing:bright yellow “lemon-aid” color, with a cloudy precip.

a.) Write the balanced net ionic equation for the equilibrium that is established upon mixing. Why does this equilibrium occur?

Observations upon addition of HNO3:1 drop cleared the solution, no color, no precipitate
a.) Account for your observations when adding HNO3(aq). Include the balanced net ionic equation and explain which way the equilibrium in (a) shifts and why in terms of Le Châtelier’s principle.b.) Which ions remain in the test tube after addition of HNO3(aq)?
Observations upon addition of HCl:Formed a solid and chunky precipitate, though not cloudy
a.) Write the balanced net ionic equation for the equilibrium that is established when HCl(aq) is added to the test tube. Why does this equilibrium occur?
Observations upon addition of NH3:Solution clears out, removing the cloudy precipitate
a.) Explain your observation when excess NH3(aq) is added to the test tube. Include a balanced net ionic equation and explain which way the equilibrium in (d) shifts as NH3 (aq) is added and why in terms of LeChâtelier’s principle.
Observations upon addition of HNO3:Remains clear, and colorless, produces a water vapor and heat.
a.) Account for your observations when adding HNO3(aq). Include a balanced net ionic equation for the reaction that occurs and explain which way the equilibrium in (e) shifts as HNO3(aq) is added and why in terms of Le Châtelier’s principle.b.) Which ions and/or solids are present in the test tube after the addition of HNO3(aq)?
Observations upon 2nd addition of NH3:No change was observed
a.) Explain your observation when excess NH3(aq) is added to the test tube. Include a balanced net ionic equation and explain which way the equilibrium in (d) shifts as NH3 (aq) is added and why in terms of LeChâtelier’s principle.
Observations upon addition of KI:Color changes to bright yellow “lemonade” color, with a cloudy precipitate
a.) Account for your observations when adding KI(aq). Include a balanced net ionic equation for the reaction that occurs and explain which way the equilibrium in (h) shifts as KI(aq) is added and why in terms of Le Châtelier’s principle.
Part III: Effect of Temperature on Equilibria
Temperature of cool CoCl2:24° C
Color of cool CoCl2 (before heating):Clear and medium red
Temperature of hot [CoCl4]2-:75° C
Color of hot [CoCl4]2- (after heating):Deep dark red, though still clear
a.) Is the reaction exothermic or endothermic? Explain your answer.

Discussion Questions

1. Based on your observations on Part I, what could be expected to happen if the solution of [Cu(NH3)4]2+ were diluted with water. Use a balanced net ionic equation to explain.

2. In the equilibrium, X3+(aq){yellow}+Y-(aq){colorless}↔[XY4]-(aq){red}, what would be indicated if heating the solution caused an intense dark red color and the solution turned yellow in an ice bath?

3. Consider the following equilibria occurring simultaneously:

Fe3+(aq){pale yellow}+SCN-(aq){colorless}↔[FeNCS]2+(aq){dark red}

Ag+(aq)+SCN-(aq)↔AgSCN(s)

2Fe3+(aq)+Sn2+(aq)↔ 2Fe2+(aq)+Sn4+(aq)

Would the color of the solution get darker or lighter if silver nitrate were added to the test tube? Explain why in terms of Le Châtelier’s Principle.

4. Would the color get darker or lighter if tin (II) chloride were added to the test tube? Explain why in terms of Le Châtelier’s Principle.

 
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