Unit+One

//__** UNIT TWO **__//
 * Lesson Date: ** Wednesday, October 16, 2013
 * Lesson Name: ** REDOX REACTIONS (OXIDATION - REDUCTION) AND ELECTROCHEMICAL CELLS
 * By: ** Kushal Gaurav

LESSON SUMMARY

 * WORK PERIOD / REVIEW DAY
 * Got points worksheet signed.
 * Unfinished points worksheets had time to finish it in class.
 * Worked on other questions from textbook and reviewed for the unit test on Firday, October 19, 2013.
 * Also, practiced nomenclature for the ISU Test tomorrow at clinic.

THINGS TO REMEMBER

 * 1) Nomenclature ISU Test tomorrow Thursday, October 17, 2013 at clinic.
 * 2) Redox Reactions (Oxidation and Reduction) and Electrochemical Cells unit test on Friday, October 19, 2013.


 * Textbook Summary **

Study Redox Reactions and Electrochemical Cells.


 * Homework **

Electrochemical Cells Review Questions Finish up points worksheet, if haven't finished yet. Practice Nomenclature ISU worksheets..


 * Extra Resource **

N/A


 * Lesson Date: ** Thursday, October 3, 2013
 * Lesson Name: ** REDOX - OXIDATION and Balancing Oxidation Numbers and REDOX REACTIONS
 * By: ** Kushal Gaurav

LESSON SUMMARY

 * Went over power point slide on oxidation - reduction again due to Ms. Wilson being away the day before.
 * Worked on example questions on the handout with the class.
 * Summary on Handout
 * Also, a new lesson was taught on balancing oxidation numbers.

**BALANCING OXIDATION NUMBERS**

 * 1) Assign oxidation to everything! (write the oxidation number above the atom).
 * 2) Identify species being oxidized and reduced.
 * 3) Balance the species oxidized between the reactant and the product side. Do the same for the species reduced.
 * 4) Determine the total loss and total gain.
 * 5) Find the Lowest Common Multiple (LCM) between the total loss and gain, use this to balance the redox participants.
 * 6) Balance everything else by inspection.


 * Textbook Summary **

N/A

Balancing Redox Reactions Worksheet
 * Homework **

Balancing Redox Reactions : []
 * Extra Resource **


 * Lesson Date: ** Wednesday, October 2, 2013
 * Lesson Name: ** Oxidation - Reduction
 * By: ** Kushal Gaurav

LESSON SUMMARY

 * Went over power point slide on oxidation - reduction.
 * Worked on example questions
 * Got a handout to fill in the blanks according to the power point.

**SUMMARY on Handout**

**WHAT IS OXIDATION-REDUCTION REACTION?**
 * ==== Oxidation - Reduction are chemical reactions where there is a loss and gain of electrons at the same time. ====
 * ====**Oxidation** - Loss of electrons====
 * ====**Reduction** - Gaining of electrons====

To make it easy to remember, there are two ways to remembering the difference between reduction and oxidation.
THESE INCLUDE: **LEO the lion says GER** or **OIL RIG.**


 * 1) ==== If an atom loses an electron = OXIDIZED ====
 * 2) If an atom gains an electron = REDUCED


 * ==== The substance gaining the electron = OXIDIZING AGENT ====
 * ==== The substance losing the electron = REDUCING AGENT ====

In redox reactions, the reaction is broken into **2 half reactions -** **//AN OXIDATION AND A REDUCTION//**

 * STEPS**


 * 1) ==== Write a balanced chemical equation. ====
 * 2) ==== Write the net ionic equation. ====
 * 3) ==== Identify what is being oxidized and what is being reduced. ====
 * 4) ==== Write the half reactions. ====

RULES FOR DETERMINING OXIDATION NUMBERS

 * 1) Elements are 0.
 * 2) A mono atomic ion is the charge of the ion.
 * 3) The algebraic sum of the oxidation numbers of a neutral compound is 0.
 * 4) The algebraic sum of the oxidation numbers in a poly atomic ion is equal t o the charge of the ion.
 * 5) Hydrogen is +1 unless in a metal hydride then its -1.
 * 6) Oxygen is usually -2 except in peroxides where it is -1. (O is +2 when with F)
 * 7) Group I metals are +1.
 * 8) Group II metals are +2.
 * 9) When second in the formula Halogens are -1.
 * 10) In combinations of nonmetals, the oxidation number of the more electronegative atom is negative, while the less electronegative atom is positive.


 * Textbook Summary **

N/A

Review handout.
 * Homework **

Today's Handout Practice Questions


 * Extra Resource **

Video we watched in class to start off the lesson. REDOX REACTIONS - CRASH COURSE CHEMISTRY : []

= = =Lesson Date: Friday, October 4th, 2013 = =[Oxidization- Reduction] =

__**//By: Nilab Ahmaddi //**__
= = =__Lesson Summary __= Ion- Electron for Balancing Redox Reaction handout


 * Started off the lesson learning how to balance redox reactions using the ion-electron method
 * The ion-electron method can be used for acidic and basic reactions
 * The following is the Ion-Election Method for balancing Redox Reactions :



__ EXAMPLE # 1 //(ACID)// __
H2S + Na +à S(s) +Na(s) H2S --> S(s) +2H+ 2 é [Na + +1é --> Na(s) ] x 2 2Na+ + 2 é --> 2 Na H2S -->S + 2H+ + 2 é
 * 2 Na + + H2S à S +Na + 2H +**

__ EXAMPLE # 1 //(BASIC)// __
MnO-4 + Ni à MnO2 + Ni 2+ 2(MnO-4 + 2 H2O + 3é --> MnO2 + 4 OH-) (Ni --> Ni2+ + 2 é) x3 2 MnO-4 + 4 H2O + 6é --> 2 MnO2 +8OH- 3 Ni --> 3 Ni 2+ + 6é
 * 2 MnO-4 + 3Ni+4H20 --> Ni2+ +2MnO2 +8OH-**

N/A
 * Textbook Summary **

The following videos help balance equations using the Ion Electron method > [] >
 * Homework **
 * complete both sides of the handout on Ion Electron Method (part b and c)
 * #|study for quiz on **WEDNESDAY!**
 * Extra Resource **
 * []


 * Lesson Date:** Friday, October 18th, 2013
 * Electrochemistry**
 * By:** Nilab Ahmaddi

=__**Summary**__=
 * Today we had our unit two - Electrocehmistry unit test
 * Answers for the test will be posted as soon as we get it back!

=__**Homework**__=
 * Read through the first 20 slides and make notes on the power point listed under the **Unit 3 Organic section** on the right side. Make sure to bring your notes to class on Monday, as we will be adding to the notes and working on problems in class.

=__//**UNIT ONE**//__=


 * Lesson Date: ** Friday, September 6th, 2013
 * Energy **
 * By: ** Vicki Lee

Under “**What is Temperature?**” on handout Under “**Temperature vs. Heat**” on handout media type="custom" key="23758256" align="center"
 * Lesson Summary **
 * Begun lesson with an activity in which we touched various classroom objects, estimating the temperatures. Afterwards, the actual temperature was measure with an #|infrared thermometer.
 * From the activity, it was observed that many of the estimates were much lower than their actual temperature values. Also, most of the objects were at about room temperature.
 * // Definitions: //
 * Temperature – the average kinetic energy of particles (i.e. how quickly they are moving); depends on speed
 * Heat – the overall thermal energy in a sample dependent on temperature, mass and type of particle (also known as “enthalpy”)
 * Watched a Eureka! video explaining, with a clever analogy, the difference between temperature and heat. []

Under “ ** Energy Changes in Chemical Systems ** ” on handout Under “**Energy Transfers**” from handout
 * // Definitions: //
 * Thermochemistry – the study of energy changes that accompany physical and chemical changes
 * // Blanks: //
 * A system that gives off heat to its surroundings is called ___exothermic___ (example: burning paper).
 * A system that absorbs heat from its surroundings is called ___endothermic___ (example: melting sugar).
 * Most energy changes are measured using a device called a ___calorimeter___ where the reaction takes place in an //isolated// compartment and the energy change can then be measured.
 * The amount of heat gained or released by a chemical system can be calculated using: **Q = mc** **∆ **** T ** = mc(T2-T1)
 * // Blanks: //
 * Heat is always transferred from an object of ___higher___ temperature to an object of ___lower___ temperature
 * Heat transfers stops only when the two objects are at the same temperature
 * Coldness is the absence of heat: an object cannot “gain coldness”, but can “gain heat”.
 * // Definitions: //
 * Conduction – transfers energy through direct contact; mostly solids
 * Convection – transfers energy through fluids (the most important method of hear transfer)
 * Radiation – transfers energy through space; no direct contact

Under “**Conductors vs. Insulators**” on handout media type="custom" key="23758270" align="center"
 * // Definitions: //
 * Conductor – a material through which heat passes through easily; will feel a lot colder than the insulator (example: metal)
 * Insulator – a material through which heat does not pass through easily (example: wood)
 * // The rest of the handout will be finished on Monday. //
 * At the end of the lesson, Ms. Wilson gave a demonstration of two ice cubes on black tiles (one a conductor and one an insulator) to show the process of heat transfer through different materials. []

n/a Study for quiz on Tuesday (safety rules, WHMIS and lab equipment) Today's Handouts 1. Energy. Please click for the PDF file. NASA: explanation of heat versus temperature: []
 * Textbook Summary **
 * Homework **
 * Extra Resource **


 * Lesson Date: ** Monday, September 9th, 2013
 * Energy (Continuation of Friday’s lesson) **
 * By: ** Vicki Lee

1.Went over unfinished slides from the "Energy" PowerPoint 2. School and TTC Photos
 * Lesson Summary **
 * // Blanks //
 * Law of Conservation of Energy equation is qgained + qlost = 0, which can then be rearranged to: **qgained = -qlost **
 * Remember: a) q = mc ∆ T b) Heat will transfer until the objects are at the same temperature
 * For the sample problem on the bottom of the page, the specific heat capacity for iron and water are 0.444 J/g°C and 4.18 J/g°C respectively. The answer to the problem is T2 = 23°C.
 * Remember that in the question, if the value of mass is given in grams (g), the unit of heat will be Joules (J). If the mass were in kilograms (kg), the unit of heat will be in kilojoules (kJ).
 * Also, you //do not// need to convert the temperature values from Kelvin to Celsius units for these questions.
 * Therefore, the “Heat Transfer Activity” has been rescheduled for tomorrow (September 10th).

n/a // Today’s Handouts //
 * Textbook Summary **
 * Homework **
 * Study for quiz on September 10th, i.e. tomorrow (safety rules, WHMIS and lab equipment)
 * Complete questions on the front side of the “Specific Heat and Heat Exchange Problems” handout according to your skill level. If a question is review for you, skip it and work on those that are on the bottom of the page.
 * 1) Appendix A: Skills Handbook, A1 Safety
 * 2) Specific Heat and Heat Exchange Problems. Please click [[file:sch4uking/Specific Heat and Heat Exchange Problems.pdf|here]] for the PDF file.
 * 3) Heat Transfer Activity. Please click [[file:sch4uking/Heat Transfer Activity.pdf|here]] for the PDF file.
 * 4) Specific Heat Capacities (Data Tables Sheet). Please click [[file:sch4uking/Specific Heat and Latent Heat Values.pdf|here]] for the PDF file.
 * Extra Resource **
 * ThinkQuest: briefly explains the Law of Conservation of energy with a pendulum example. Some physics review is included at the bottom, if it interests you. []
 * AllThink: gives a step-by-step tutorial on solving heat questions. []


 * Lesson Date: ** Tuesday, September 10th, 2013
 * Heat Transfer Activity **
 * By: ** Vicki Lee

1. Safety and Lab Equipment Quiz 2. Heat Transfer Activity
 * Lesson Summary **
 * We wrote the quiz during the first half of the class.
 * The answers to the quiz questions will posted here when we get them back.
 * Please click [[file:sch4uking/Heat Transfer Activity.pdf|here]] for the PDF file.
 * Worked through the activity in pairs after writing the quiz.
 * The purpose of the activity was to find the identity of an unknown metal by calculating its specific heat capacity with the data measured by a calorimeter. It also incorporates what we’ve learned, so far, about heat transfer conceptions and calculations.
 * Boiling water and a piece of unknown metal was placed in a calorimeter (made with two Styrofoam cups stacked together and half a cup as a lid) and the initial temperature was recorded.
 * During the heat exchange process, we worked on the Calorimeter questions on the reverse of yesterday’s handout while we waited. At the end of the period, the final temperature was recorded.
 * You can watch a video briefly explaining the activity here: https://www.youtube.com/watch?v=lC71SVTGYO4

media type="custom" key="23786824" align="center"


 * Tomorrow, we will be doing the “**Phase Change Activity**” on the back of the “**Heat Transfer Activity**” handout. Please click [[file:sch4uking/calorimetry and phase questions.pdf|here]] for the PDF file.

n/a 1. Finish calculations on the Heat Transfer Activity sheet 2. Complete questions under the “ ** Calorimetry Questions ** ” heading on the “ ** Specific Heat and Heat Exchange Problems ** ” handout from yesterday. // No handouts today. // YouTube video explaining how to solve a calorimeter question, similar the one from today’s “Heat Exchange Activity” (only the unknown substance is a liquid) in a step-by-step tutorial. [] media type="custom" key="23786812" align="center"
 * Textbook Summary **
 * Homework **
 * For questions involving calorimeters, remember that there must be a separate term for the calorimeter in the equations (i.e., it’s own m, c and ∆ T values, either on the qgained or qlost side).
 * Extra Resource **

=**Lesson date:** Wednesday, September 9th 2013= =**Phase Change Activity**= =**By:** Samantha Karvanis=

**Lesson Summary:**
-Performed lab activity in partners, to investigate heat changes that accompany phase changes -Around 50g of ice was placed into a beaker that was then put onto a hot plate to melt the ice -The temperature of the ice was taken every minute as it was being heated, and the temperature was then recorded in a data table -The temperature was recorded up to 4 minutes after the water was boiling - The beaker was then removed from the hot plate -The data that was recorded during the activity was then graphed -The graph illustrated the rate of temperature increase as it changes state -The total amount of heat absorbed by the water was calculated at the end of the activity
 * 1)** Phase Change Activity

Homework:
1) Phase Diagram Questions - Question #'s 1-7 on the back of the "Specific Heat and Heat Exchange Problems" worksheet -Below is an attached video that deals with phase diagrams. It discusses topics that we already learned like specific heat capacity, and also goes into specific latent heat. It also demonstrates how to do problems involving latent heat: media type="custom" key="23798252" align="center"

- Below is another attached video on phase change diagrams:

media type="custom" key="23798328" align="center"

=Lesson Date: Thursday September 12, 2013=

Summary by Anne Clayton
- Following our phase change activity, we learned that the total amount of heat absorbed by the water could not be calculated by using only the equation q=mc ΔT, since the c value for ice (2.10) is different than that of liquid water (4.18). - We also learned why plateaus occurred on our graphs - these were the times when a phase change was occurring.

- A phase change occurs when a substance changes state - During a phase change the temperature of the substance does not change. The energy goes towards weakening the forces between molecules. This appears as a plateau on a phase change diagram. - In order to calculate the amount of heat absorbed (or released) during a phase change, **latent heat values** must be used. There are two values; that for **fusion** (melting) and that for **vaporization** (boiling). - For questions involving freezing and condensation, use the negative values for fusion and vaporization.

Above is a rough replication of the example of a phase change diagram shown in class. The plateaus show the areas where a phase change is occurring, also known as areas where you must use the latent heat value to find out how much energy was required to change states **(q=lm)** where l is the latent heat value and m is the mass. The areas of increase show the substance heating up also known as areas where **q=mc** Δ**T** can be used.

If you would like a further explanation of phase changes. here is a very informative video which explains it thoroughly.

**REMEMBER!!** QUIZ NEXT WEEK!!

__** Lesson Summary: **__

 * Today we learned more about enthalpy, which is the amount of heat produced or used up during a chemical reaction and can be represented using the symbol ΔH
 * ΔH has the unit of KJ/mol and is also equal to the differences in enthalpies of reactants and products.
 * Remember that the ΔH for an exothermic reaction will be negative as heat it being __released__ and for an endothermic reaction ΔH will be positive as heat is being __absorbed.__
 * We have a couple of ways to calculate ΔH they include by:
 * 1) Using the values of q and the number of moles.
 * 2) Using Hess' Law
 * 3) Using the Standard Heats of Formation
 * 4) Using Bond Energies
 * 5) Potential Energy Diagrams

__The Standard Enthalpy of Formation__
> ΔH from our Enthalpies of Combustion handout for the rest of the compounds that remain.
 * The Standard enthalpy of formation (//H//fO ) is the the energy that is associated with creating a substance from its elements.
 * Remember that elements have their standard enthalpy of formation of zeros. The values for compounds can be found on the Appendix C: Specific Heat Capacities handout that we got a few classes ago.
 * Also note that when you solve for ΔH, begin with a balanced chemical equation and include states!
 * Normally in Grade 11 we learned to balance the chemical equation, but here we keep the co-efficient for the fuel as one, so fraction are okay for this case.
 * So, C8 H18 (l) + 25/2 O2 (g) = 9 H2O(g) + 8 CO2 (g) is okay.
 * We learned that:ΔH = Σ ΔHf products - Σ ΔHf reactants
 * To use this equation, we just omit the elements and plug in the values for

__Homework__

 * Finish __all__ the questions from Calorimetry, Phase Diagram and Specific Heat and Heat exchange worksheet. They will be checked on Monday!
 * Finish the Standard Heats of Formation worksheet that is posted on wiki space. A direct link can be found here
 * Remember that our quiz is on Friday! It will cover everything we have learned in the energy unit.

__Resources__

 * [|Here] is a sample worked problem on using the standard heats of formation.
 * Here is a video that goes more in depth on Enthalpy. In the video's description you can navigate through the sections which interest you the most.

=Date of Lesson: September 16, 2013= =Bond Energies and Representing Energy Changes Graphically=

By: Anne Clayton
__**Indicating ΔH for a Reaction **__ - ΔH can be written after the reaction, as follows: H2 (g) + ½ O2 (g) à  H2O (l) ΔH= -285 Kj/mol of H2 - it can also be written where the amount of energy is written as a term in the equation as follows: H2 (g) + ½ O2 (g) à <span style="font-family: Arial,sans-serif; font-size: 12pt;"> H2O (l) + 285 kJ

__**Calculating**__ ** __ΔH__ **__ **Using Bond Energies** __ - Bond breaking is an **endothermic** process and bond forming is an **exothermic** process - therefore enthalpy of a reaction can be calculated using bond energies ΔH = Σ bonds broken - Σ bonds formed **Before summing, bond energies must be multiplied by the number of that bond present AND the coefficient from the balanced chemical equation**
 * - The structures of each species must be drawn to determine the number and type of bonds present **


 * Here is a picture of a sample problem we did in class, for the combustion of acetone.**


 * __ Representing Energy Changes Graphically __**

- Changes in energy in a chemical system can be represented graphically using a **Potential Energy Diagram** - These graphs plot potential energy vs reaction progress - ΔH is calculated by determining the difference in potential energy between the reactants and the products

-Potential Energy Diagrams never go to 0 because all chemicals have some type of stored energy - In an endothermic reaction, the reactants are lower than the products - in an exothermic reaction, the reactants are higher than the products

[|Here] is an informative video on how to complete bond energy problems. [|Here]is a video with additional information on potential energy diagrams.

Remember, there is a quiz on Friday about everything we have learned!!

__Source__s [] []

=Date of Lesson: Tuesday, September 17 2013=

By: Anne Clayton
- ΔH for a reaction can be calculated algebraically for a target reaction by using various related reactions (reference reactions) and Hess' Law. - If a reaction must be **reversed,** the sign of the ΔH must be **changed from positive to negative, or from negative to positive** - While adjusting molar ratios by multiplying or dividing **all coefficients as well as the value of** ** ΔH must be multiplied or divided by the same factor **

Sample Problem
<span style="font-family: Arial,sans-serif; font-size: 10pt;">Determine ΔH for C2H6 (s) à <span style="font-family: Arial,sans-serif; font-size: 10pt;">C2H4 (g) + H2 (g) using…
 * 1)** <span style="font-family: Arial,sans-serif; font-size: 10pt;">2 C2H6 (s) + 7 O2 (g) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> 4 CO2 (g) + 6 H2O (l) ΔH = -3 118 kJ/mol
 * <span style="font-family: Arial,sans-serif; font-size: 10pt;">2) **<span style="font-family: Arial,sans-serif; font-size: 10pt;">C2H4 (g) + 3 O2 (g) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> 2 CO2 (g) + 2 H2O (l) ΔH = -1 411 kJ/mol
 * <span style="font-family: Arial,sans-serif; font-size: 10pt;">3) **<span style="font-family: Arial,sans-serif; font-size: 10pt;">H2 (g) + ½ O2 (g) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> H2O (l) ΔH = -236 kJ/mol

<span style="font-family: Arial,sans-serif; font-size: 10pt;">In order to get 1 C2H6 equation **1** must be divided by 2 <span style="font-family: Arial,sans-serif; font-size: 10pt;">In order to get C2H4 on the products side, equation **2** must be flipped along with the sign on the ΔH value <span style="font-family: Arial,sans-serif; font-size: 10pt;">In order to get H2 on the products side, equation **3** must be flipped along with the sign on the ΔH value


 * 1)** <span style="font-family: Arial,sans-serif; font-size: 10pt;">C2H6 (s) + 3/2 O2 (g) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> 2 CO2 (g) + 3 H2O (l) ΔH = -3 118 kJ/mol
 * <span style="font-family: Arial,sans-serif; font-size: 10pt;">2) **<span style="font-family: Arial,sans-serif; font-size: 10pt;">2 CO2 (g) + 2 H2O (l) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> C2H4 (g) + 3 O2 (g) ΔH = 1 411 kJ/mol
 * <span style="font-family: Arial,sans-serif; font-size: 10pt;">3) **<span style="font-family: Arial,sans-serif; font-size: 10pt;">H2O (l) à <span style="font-family: Arial,sans-serif; font-size: 10pt;"> H2 (g) + ½ O2 (g) ΔH = 236 kJ/mol

<span style="font-family: Arial,sans-serif; font-size: 10pt;">All species that appear on both sides must be crossed out to leave…

<span style="font-family: Arial,sans-serif; font-size: 10pt;">C2H6 (s) à <span style="font-family: Arial,sans-serif; font-size: 10pt;">C2H4 (g) + H2 (g) add the ΔH values 236+1411-3118= 88 kJ/mol

<span style="font-family: Arial,sans-serif; font-size: 10pt;">Therefore the ΔH of the reaction is 88 kJ/mol.

[|Here] is a video demonstrating how to use Hess' Law. [|Here] is a PowerPoint from a school demonstrating the use of Hess' Law.

Remember there is a quiz on Friday on everything we have learned including today's lesson.

Date of lesson: September 18th 2013
=Rates of Reaction=

__Rates of Reactions__
- the rate of a chemical reaction is the speed at which a chemical reaction occurs (e.g. speed at which the reactants are used or products are produced) -rate can be measured by the: 1)change in mass of reactants or products 2) change in pH 3) change in conductivity (ion production) 4) change in colour (intensity of colour) 5) change in temperature 6) production of gas


 * if rate is large then it has a small amount of time in which the reaction is formed
 * when taking temperature of gas it must be in kelvins

__Calculating Average Rate of Reactions__
- average rate of reaction is the change in concentration in a given time period -mathematically, rate= Δ[C] / Δ**t** rate = ([C]final -[C] initial) / (**t** final - **t** initial) [C] is concentration in mol/L and **t** is time (usually seconds) The unit for rate is mol/L s - rate is expressed in terms of t reactant is consumed - reactants are used and products are used based on their molar ratios

Here is a video that demonstrates how to do rates of reactions problems. It is done a bit differently then the sample problem we did in class, but it gives you the general idea. media type="custom" key="23870208"

Date of lesson: Thursday, September 19th 2013
=Heat of Combustion of Magnesium Lab=

Heat of Combustion of Magnesium Lab
- Performed the lab in partners - Carried out two reactions for which the heat change involved was calculated - We then used these values to calculate the heat of combustion of magnesium

Reaction A: Mg(s) + 2 HCl(aq) ---> MgCl2(aq) + H2(g)
-placed 100 mL of 1.0 M hydrochloric acid into a styrofoam cup -the temperature of the solution was recorded - the mass of 1.0g of magnesium ribbon was found - the magnesium was then added to the acid - the solution was stirred carefully and the highest temperature of the solution was recorded

Calculate:
-the number of moles of Mg used in the experiment -the heat gained by the solution (use the specific heat capacity for water) -the heat of reaction per mole of Mg

Reaction B: MgO(s) + 2 HCl(aq) ---> MgCl2(aq) + H2O(g)
- dumped out the solution from reaction A - placed 100 mL of 1.0 M hydrochloric acid into a styrofoam cup - the temperature of the solution was recorded - the mass of 1.0g magnesium oxide was found - the magnesium oxide was added to the acid - the solution was stirred carefully and the highest temperature of the solution was recorded

Calculate:
-the number of moles of MgO used in the experiment -the heat gained by the solution (use the specific heat capacity for water) -the heat of reaction per mole of MgO

Follow Up Questions:
1)Calculate the heat of combustion of magnesium using your experimental data 2)Calculate the heat of combustion of magnesium using heats of formation 3)Calculate your percent difference between your answers from 1 and 2

Here is a video that is similar to the lab we did:

[]

** Summary: **

 * ==== Today we used the majority of our class period working on our first Energy Quiz. The quiz answers will be posted shortly after we get them back. ====
 * At the end of today's class we also obtained our first ISU. It is a general review of the nomenclature that we all learned in grade 11. Remember to continue working on this during your own time as this is an ISU.
 * In case you forgot to grab one or misplaced the handout, a direct link to the worksheets can be found here.
 * Once you are done everything, make sure you have all the correct answers by checking with the answers here.

**Homework:**

 * Start working on the Nomenclature Independent Study. There will probably be a test on it in the near future.
 * Also, as a heads up our Energy Test, as of now is scheduled for October 1, 2013. So may want to start doing some review.

**Resources:**

 * In cause you are still having some troubles with Energy, here are some extra practice questions.
 * Here is a great video that goes over the basic knowledge of chemical nomenclature.

__**Quiz Answers**__



Summary:

 * ====Factors that affect the rate of reaction====
 * 1) Temperature
 * 2) Surface Area
 * 3) Concentration
 * 4) Catalyst
 * 5) Chemical Nature of Reactant
 * 6) State of the reactant

Collision Theory:

 * During a reaction the reactant particles move around and collide with each other
 * Some collisions will cause the bonds in the reactants to break and new bonds will form making new molecules
 * Those collision are called effective collisions
 * The rate of reaction increases by increasing the number of total collisions and increasing the number of effective collision by decreasing the activation energy

Temperature:

 * Temperature is the average kinetic energy in molecules
 * By increasing the temperature it will increase the total number of collisions
 * Which will increase the effective collisions and increase the rate of reaction

Surface Area:

 * Increasing the amount of particles that could react
 * Which will increase the number of total collisions and that will increase the number of effective collisions
 * Which will increase the rate of reaction

Concentration:

 * The greater number of particles the greater the number of total collision
 * The total number of collision increases so the effective collision increases as well
 * Which will increase the rate of reaction

Catalyst:

 * A catalyst works by providing an alternative pathway for the reaction which has a lower activation energy
 * Lowering the activation energy will only increase the number of effective collisions
 * The number of total collisions will NOT be affected

Chemical Nature of Reactants:

 * The type of reactant determines the activation energy required
 * The higher the activation energy, the slower the rate of reaction
 * Endothermic reactions are much slower than exothermic reactions because they tend to have higher activation energies

State or Phase of the Reactants:

 * Reactions where all the reactants are in the same state (homogeneous reaction) will occur at a faster rate than reactions where reactants are in different states (heterogeneous reaction)
 * This is because the reactants will have a greater opportunity of colliding
 * The state of the reactants will also affect the rate
 * The relative rates are
 * Gases are the fastest, liquids/solution are fast, solids are slow
 * Stirring increases the rate of reaction

Homework:

 * Continue to work on the nomenclature ISU
 * The energy and rates test has been scheduled for Tuesday October 1, 2013, so begin to study
 * Start to design the rate of reaction lab

Extra Help:
>
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**Date of Lesson: September 24, 2013**

Rate Law of Reactions:
ra[X]^m [Y]^n (where a represents proportional)
 * The relationship must be determined empirically (must analyze experiment data)
 * The rate of law for a reaction describes the relationship between rate (r) and the product of the initial concentrations of the reactants raised to some exponential values
 * Given: aX + bY --> products
 * The exponents m and n describe the relationship between rate and initial concentration and must be determined experimentally
 * The value of m and n may be a whole number, zero or a fraction and do not have to equal the coefficient from the balance chemical equation
 * A Rate Law Equation for the reaction is written as r = k [X]^m [Y]^n
 * The value for the rate constant must also be determined experimentally

Orders of Reaction:

 * The exponents in the rate law equation are called the orders of reactions
 * The sum of all the individual orders (exponents) is referred to as the overall order of reaction
 * Example: Given 2X + 2Y + 3Z --> products
 * The order of reaction X is 1, Y is 2 and Z is zero (which has no affect on the reaction) so the overall order is 3
 * In 1st order, if concentration is doubled the rate is doubled (2^1)
 * In 2nd order, if concentration is doubled rate is quadrupled (2^2)
 * In 3rd order, if concentration is doubled rate is increased by a factor of 8 (2^3)
 * In 0 order, if concentration is doubled, there is no change in reaction

Rate Sample Problem:
The was data below was collected from reactions

A(aq) + 2 B(aq) + 3 C(aq) --> E(s) + F(aq)

Determine the rate law equation, including the value of k

Rate || To start with A you must find which two #|experiments have the same B and C values but different A values, so experiment 2 and 3
 * A || B || C || Initial
 * 0.10 || 0.10 || 0.10 || 4.0 ||
 * 0.10 || 0.20 || 0.10 || 16.0 ||
 * 0.40 || 0.20 || 0.10 || 32.0 ||
 * 0.10 || 0.10 || 0.30 || 12.0 ||

For A, compare 2 and 3 To get from 0.10 in experiment 2 to 0.40 in experiment 3 you must multiply by 4 so [A] x 4 and to get from 16.0 to 32.0 you must multiply by 2 so r x 2. Now you must get both multipliers to be in exponential form with the same base, so the first one is already 4 and the second one is 2 so to get to the same base you would change 2 to 4^1/2 then your exponent will be the exponent used in the equation so A = 1/2

For B compare 1 and 2 [B] x 2 r x 4 = 2^2 --> 2nd order

For C compare 1 and 4 [C] x 3 r x 3 = 3^1 --> 1st order

Then you put any experiment and plug in the values using the exponents you found, so I will use experiment 2

r = k[A]^1/2 [B]^2 [C] 16.0 = k[0.10]^1/2 [0.20]^2 [0.10} k = 1.3 x 10^7

Homework:

 * Finish the rate of law reaction worksheet
 * Study for upcoming test on October 1st
 * Work on your design for the rate of reaction lab
 * There will be an ISP for class tomorrow so pick up the handout from the classroom and you can work in the classroom

Extra Help:

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**Date of Lesson: September 25, 2013**

Homework

 * Work on the sheet assigned
 * Study for upcoming unit test on Oct 1st

**Lesson Summary**
a) Activation Energy (Ea), the minimum amount of potential energy required for a reaction to occur b) Ea for the reverse reaction c) <span style="font-family: Arial,sans-serif; font-size: 13px; line-height: 1.5;"> Δ H Activated Complex: Reaction intermediates (collision state)

If the reaction is heater it would translate the whole function up.

Potential Energy
• The lower the activation energy, the greater the rate of reaction.

• In general, exothermic reactions tend to have greater rates than endothermic reactions.

• Many exothermic reactions are __ self-sustaining __ *Endothermic can never be self-sustaining because they use energy


 * <span style="font-family: Arial,sans-serif; line-height: 1.5;">ΔH | > E - Produce enough energy to overcome the activation energy.
 * Catalysts increase rate providing an alternative pathway with a lower activation energy.

Maxwell - Boltzman Distribution

 * Bigger curve for concentration.
 * Temperature shifts curve to the left.




 * To increase the rate of reaction, the fraction of particles with the activation energy or greater must increase.


 * __ Increase Temperature __**
 * Lower temperature will be T 1, and higher temperature will be T 2

__ With a Catalyst __

The distribution of particles kinetic energies at a fixed temperature (at the top of the curve)


 * Ea 1 the higher activation energy of the uncatalyzed reaction
 * Ea 2 the lower activation energy of the catalyzed reaction

Example

 * Consider the production of NO 2 from NO and O 2

2 NO + O 2 -> 2NO 2

This reaction consists of elementary steps.

1) 2 NO -> N2O 2 (Fast, molecularity 2) * Both biomolecularity

2) __ N ____ 2 ____ O ____ 2 ____ + O ____ 2 ____ -> 2NO ____ 2 __ (Slow, molecularity 2)

2NO + O 2 -> 2NO 2

N 2 O 2 is a reaction intermediate which is short lived and difficult to isolate.

Reaction Mechanisms

 * A reaction mechanism is step or series that makes up a reaction.
 * The step ina a reaction mechanism are referred to as **elementary steps.**
 * **Molecularity:** Refers to the number of reactant molecules involved in an elementary step.
 * **Unimolecular:** Steps involving 1 molecule.
 * **Bimolecular:** Steps involing 2 molecules.
 * **Termolecular:** Steps involving 3+ molecules.
 * Lower molecular steps tend to be faster than ones with more molecules.
 * **The rate determining step** is the slowest step in the reaction mechanism and will have the largest activation energy.

Using Reaction Mechanism to Determine the Rate law Equation

 * The rate law expression can be determined from the reaction mechanism.
 * Rate determining step is the elementary step which determines the rate law expression.
 * Rate of reaction is proportional to the concentrations of the reactants in the rate determining step raised to their molar coefficients.
 * Any intermediates that appear in this equation, you must use the other steps to eliminate it.

Homework

 * Complete the back of the Reaction Mechanism sheet
 * Study for test that is on Tuesday October 1

Extra Resources
[] Help with reaction mechanism problems

Summary

 * Today we had the full period to catch up and finish up all the worksheets we got in this unit.
 * Make sure you guys all have the Energy handouts we got throughout this unit. If you missed one, check the main Wiki to print yourself a copy. Click here to go to the main wiki.
 * All the concepts that we learned throughout this unit are important, and will show up on the test in one form or another. So make sure to do all the practice questions and complete all the worksheets.
 * Remember our unit test is on **Tuesday October 1, 2013!** !
 * When you feel like you are ready, give this practice Energy test a try, to see how well you do. The solutions for this practice test are posted just below it.
 * The remainder of the period was used to continue working on our Rates of Reaction Lab Design.
 * We were also given back our first Energy Quiz that we all wrote last week.
 * Please check my other wiki post (on September 20 2013) for the correct answers for the quiz.
 * Also we were informed that Ms. Wilson won't be here on Monday, and that most likely we will be getting an ISP. Keep an eye on your wiki mails, as Ms. Wilson said that she would be sending us more info on that, in regards to what we are supposed to be doing during that time.

Homework

 * Study for the Energy Unit Test that is next Tuesday!!
 * Finish working on the Rates of Reaction Lab Design.

Extra Resources

 * If you feel like you need some extra practice this site for some practice questions and some sample problems regarding Calorimetry questions.
 * For more practice on Hess' Law check out this website. This site also has some great examples with solutions.
 * Click here to review how to do some of the phase change questions.

Homework

 * STUDY FOR OUR TEST TOMORROW!
 * There was a handout on the wall outside of our classroom that has a [[image:Test 1.jpg]]






 * Review from our text book.
 * Some good practise questions from that booklet are #7-17, 19, 20, 23, 25-28, 30-32, 36a, 37, 39-41, 44, 45, 50-53 (and for number 53 enthalpy diagram is a potential energy diagram) 65 and 69.
 * There was also a self-assessment, but you do NOT have to worry about numbers 9, 15, 18, 19-21, 24 and 25.
 * Good luck everyone!

The answers will be posted once we get them back!
=**Date: October 7, 2013 - Work Period + Homework Take-Up**=


 * Class Summary:**

Today we had a work period where we took up questions and were given time to finish up Part A, B and C of the Redox reaction balancing worksheets to prepare for our quiz on Wednesday.


 * Homework:**


 * Part A, B and C of the Redox Balancing Practice
 * Study and Practice for Quiz

[] [] http://pages.towson.edu/ladon/redoxprac.htm
 * Extra Links/Practice for Redox Balancing **

Benett Szuba

** Date: October 8, 2013 - ISP **

 * Class Summary:**


 * Today was an ISP and we received a short reading on different types of cells and an introduction to Electrochemistry.**


 * Homework:**
 * **Finish reading**
 * **Study and practice for balancing quiz**

Date: October 9, 2013 - Quiz Day (Answers will be posted when available)

 * Class Summary:**


 * Today we did a Redox reaction balancing quiz. Later in the class we received our text books and an ISU sheet that will be explained tomorrow.**


 * Homework:**
 * **Read section 9.3 (Page 618 - 622)**
 * **Finish reading from ISP**


 * 9.3 - This section discussed predicting Redox reactions, the strengths of different Oxidizing and Reducing Agents and it also contained some practice problems.**

BY: RAMMIYA JEGANATHAN

**DATE: October 10th 2013**

 * Class Summary:**


 * Today we learned about Electrochemistry, types of electrochemical cells, electron flow in a cell.**

__ ** ELECTROCHEMISTRY ** __

The study of chemical reactions that convert chemical energy into electrical energy



//** This is a galvanic cell **//

__** ELECTROCHEMICAL CELLS **__
 * A device that continually converts chemical energy to electrical energy

__**PARTS**__


 * 1) **//Electrodes (2)//**: Solid electrical conductors connected in some manner where the redox reaction occurs. Allows for metallic conduction (flow of electrons)
 * 2) //**Electrolyte:**// An aqueous solution (or paste) that the electrodes are immersed in. Allows for electrolytic conduction (flow of ions)
 * 3) //**Wires:**// To connect the electrodes to the electrons can flow.

For the cell to function properly it must be connected externally for example electrodes must be connected and internally for the flow of ions. If the electrodes are in two different electrolytes. A salt bridge is used to link the two half cells.

//**Salt Bridges**//: Contains a salt solution and allows the ions to move to the other container in order to maintain the electrical neutrality of the solutions. (U-tube)

Positive ions always move to the cathode and the negative ions go to the anode.
 * Remember:**

__**ELECTRON FLOW IN A CELL**__


 * In every electrochemical cell oxidation occurs at the ANODE and reduction occurs at the CATHODE.
 * As a result, electrons flow from the ANODE to the CATHODE in every cell
 * The force pushing the electrons around the cell in the potential difference or electromotive force that exsists between the two electrodes also known as voltage
 * The rate at which electrons flow is the current

__**TYPES OF ELECTROCHEMICAL CELLS**__
 * There are two main type of cells, galvanic or voltaic and electrolytic

//__TYPES OF ELECTROCHEMICAL CELLS__// to generate current || -use energy released from a non-spontaneous redox reaction to generate current || __**REDOX TABLE**__
 * **Galvanic or Voltaic** || **Electrolytic** ||
 * - use energy released from a spontaneous redox reaction
 * -oxidation occurs at the ANODE || - oxidation occurs at the ANODE ||
 * - Reduction occurs at the CATHODE || - Reduction occurs at the CATHODE ||
 * - ANODE = Negative electrode || - ANODE = Positive electrode ||
 * - CATHODE= Positive electrode || - CATHODE= Negative electrode ||
 * - Electrons flow from the anode to the cathode; negative to postive || - Electrons flow from the anode to the cathode; positive to negative ||
 * - No power supply is required || - A power supply is required ||
 * EX. Discharging a Battery || EX. Recharging a battery ||


 * In a redox table, all half reactions are shown as reductions so to determine the oxidation half cell reaction REVERSE the reaction and change the sign of the E cell value
 * In order for the redox reaction to be spontaneous the sum of the cell potentials must be POSITIVE

eg. Will the oxidation of Cr to the Cr 3+ ion by Mg 2+ be spontaneous?

E o v ANODE : Oxidation : Cr ---> Cr 3+ + 3e- 0.74 V CATHODE: Reduction : Mg 2+ ---> Mg (s) __-2.37 V__ -1.63 V Therefore, this reaction is NOT spontaenous, you need 1.63 V to make it go.

__**STANDARD CELL NOTATION:**__

Using the previous sample question we will write out a standard cell notation



//**- Characteristics of Electrochemical cells :**// Answer pg 641 #3-7 and pg 670 #1,2 Draw a diagram of a glavanic cell, label all parts, indicated the funtion of each part, show the electron flow, show the ion flow and give the standard cell notation for your cell //**- Predicting Redox Reactions and Standard Cell Potentials:**// Complete the Electrochemical Cells Worksheet //**- Electrolysis and Corrosion**// Summarize pgs 663 - 670 Summarize pgs 658 - 661
 * Homework:**
 * **Electrochemistry Layered Curriculum**
 * PICK ANY THREE IN THE LIST**
 * Unit test on FRIDAY
 * Nomenclature ISU Test On October 17th, 2013


 * LESSON DATE: October 15th 2013 By: Rammiya Jeganathan**
 * Homework:**
 * Review for test, finish all worksheets etc.**

We did our nomenclature test today We had a reviewing session for tomorrow's test today.
 * LESSON DATE: October 17th 2013 By: Rammiya Jeganathan **

Homework**:**
 * STUDY FOR THE UNIT TEST THAT IS TOMORROW !!!!**