Solutions & Stoichiometry

It's finally here... Solutions + stoichometry!

Okay, so you might not be so excited about this. But hopefully this blog

can clear up any confusions you have about blending what we already

know about stoichometry with this unit's topic: solutions.

• Today, we started class by turning in our Molarity Stiumlations packets, which we're on schedule to go over tomorrow. If you need one of these, they're in the Unit 4 Handouts folder!

• Then, we started on the torture part of this unit: Solution stoich. 

Here's a brief overview of the notes we took and the problems we did!

Disclaimer: Ms. Friedmann's notes, which she wrote in class today, are in the Unit 4 Notes folder 

under  'Notes on Solution Stoichiometry'. Go check these out, as well.

First, we redefined stoichometry: looking at one "thing" in a reaction (either a reactant or a product) and figuring out how much of another "thing" in the same reaction will be used up or produced based on the measurement of "thing" 1. 

This definition might seem a little extensive, but to remind everyone...

A stoichiometry set-up looks just like this! 

Now, we all remember doing (a little bit too many of) these problems!

So, where do solutions come in?

Well, to solve stoichometry problems involving Molarity, we need new algorithms.

Our old algorithm looked like this:

We've used this algorithm to solve all of our stoichiometry problems to date. But now,
with solutions, we need a new algorithm involving molarity.
The new algorithm:

Obviously, the new parts of this algorithm are:

• Calculations using the volume to find other measurements. Use the molarity to convert from volume to moles, and vice versa.
• Calculations using the molarity to find other measurements. Use the volume to convert from molarity to moles, and vice versa.

Let's do a problem:

This is the first problem on the worksheet entitled 'solution stoich', located in the Unit 4 Handouts folder.

CaCl2 (aq) + 2 AgNO3 (aq) -----> Ca(NO3)2 (aq) + 2 AgCl (s)

What volume of 0.150 M AgNO3 is needed to react with 45.0 mL of 1.50 M CaCl2?

What is "thing 1"? CaCl2, because we are given its molarity and volume.

What is "thing 2"? AgNO3, because we are asked to calculate the volume of it needed to react.

There are two ways to start this problem, either with the molarity or the volume of CaCl2. We chose to start with the volume of CaCl2.

Since the volume is in mL, we must convert to L just by shifting the decimal 3 places to the left.

Knowing our first conversion factor must be 1.50 moles CaCl2/1L, because molarity is moles of solute over liters of solvent, we can begin to set this problem up.

0.045 L * 1.50 moles CaCl2/1L

Now, we need to convert from moles of CaCl2 to moles of AgNO3. To do that, we look at the balanced equation. There are 2 moles of AgNO3 for every 1 mole of CaCl2. So...

0.045 L * 1.50 moles CaCl2/1L * 2 moles AgNO3/1 mole CaCl2

Next, we need to get from moles to volume of AgNO3. So, knowing the molarity of AgNO3 and using the corresponding conversion factor...

0.045 L * 1.50 moles CaCl2/1L * 2 moles AgNO3/1 mole CaCl2 * 1L/0.150 moles AgNO3 =

0.90 L of AgNO3!

And that's it! The rest of this sheet, along with the two others recieved in class today, are for homework. Good luck, and feel free to ask questions in the comments section!

Here's an awesome Khan Academy video on solution stoichiometry if you're still confused:

Next scribe is Mary L!

Solubility Lab

Solubility Lab

Class on October 31 (Halloween)

Written by: Jordan C.

After the class had a short discussion of Shakespearean iambic pentameter, we got down to business.

Remember that soluble means that the ionic bond is weaker than the water bond and therefore remains separated.

Going over the double replacement reactions and solubility lab:

• Pre-lab Questions: The class did not go over these because they were not assigned.
• Data: The data did not have to corrected because the rules as to the solubility and insolubility with cations and anions is more important. 
• Post-lab Questions: Ms. Friedmann gave an example of each of the questions and demonstrated how to find the answer to any of the class's requests. 

Question 1:

• Although the 2 possibilities for the possible precipitate were required, Ms. Friedmann saw this as an opportunity to demonstrate a balanced molecular equation for:

Combination of: Ba(NO3)2 and Na2SO4. 

Ba(NO3)2(aq) + Na2SO4(aq) ---> BaSO4(s) + 2 NaNO3(aq)

Ba(NO3)2(aq) + Na2SO4(aq) ---> BaSO4(aq) + 2 NaNO3(s)

Question 2: 

• The demonstration included the combination of Ba(NO3)2 + Na2CO3. She also showed the net ionic and ionic equations for practice. 

Molecular: 

Ba(NO3)2(aq) + Na2SO4(aq) ---> BaSO4(s) + 2 NaNO3(aq)

Complete ionic:

Ba^+2(aq) + 2 NO3^-(aq) + 2 Na+(aq) + SO4^-2(aq) --->BaSO4(s) + 2 Na^+2(aq) + 2 NO3

Net ionic:

Ba^+2(aq) + SO4^-2(aq) ---> BaSO4(s)

Question 3:

a) SO4^-2, halides (ions of halogens ) Cl^-, I^-

b) silver cations Ag^+1 and barium cations Ba^+2

Question 4: 

a) CO3^-2 carbonates, OH^- hydroxides, PO4^-3 phosphates

b) potassium cations K^+, sodium ions Na^+1 (THESE ARE ALKALI METAL IONS)

Question 5:

a) Bromine salts would be soluble because it is in the halogen family, and therefore is closely related with the other soluble halogens.

Question 6:

This is one of the most important aspects of the lab. All the information leads up to discovering the rules of solubility with cations and anions. 

Ionic Compounds:

• Carbonates: insoluble except alkali metals (sodium, potassium)
• Halides: soluble except silver
• Hydroxides: insoluble except alkali cations (K^+1)
• Nitrates: soluble except none :)
• Phosphates: insoluble except alkali cations
• Sulfates: soluble except barium (Ba^+2)
• Alkali Metal salts: soluble except none :)
• Ammonium salts: soluble except none :)

As class came to a close, we had the opportunity to make bubbles and bounce them off gloves. It was suuuuper fun. 

Thursday, 10/31 Homework

1) Complete the "Solubility Rules and Practice" worksheet (in the Unit 4 Handouts folder).  Due tomorrow.

2) WebAssign 4.2 -- Precipitation Reactions.  Due Sunday night, 11/3, 11:59 pm.

Next Post will be by...

Ambreen A.!

Double Replacement Reactions and Solubility Lab

Scribe: Suvd D.

Tuesday, October 30th, 2013

Homework:
Complete the post-lab questions for the "Precipitates Lab", posted in the Unit 4 Labs folder.
Note that for question #2 you only have to pick eight of the precipitation reactions and write only molecular equations for them...we will do net ionic equations for them tomorrow.  

Notes on how to write a molecular/complete ionic/net ionic equation are posted in the Unit 4 Notes folder.

These are different examples of equations that Ms. Friedmann gave us in class.

First of all, Ms. Friedmann explained the video on Juliette's blog. She said that it was an exothermic process which means that the process releases heat. The supersaturated solution has lots of free floating ions that if you put in a little bit of crystal , the other ions start sticking onto the crystal. This is what it looks like after:

Here's Juliette's video:

Fun but Dangerous...
You can make your own supersaturated solution!
All you need is baking soda, vinegar and heat!
But...it can burn your lungs so be careful.
The formula: NaHCO3 (baking soda) + HCH3COO (vinegar)--> NaCH3COO (aq) +H2CO3

Other announcements:
This quarter, Ms. Friedmann wants us to scribe with a new label "Q2 Name". She also wants us to comment on each other's blogs. Lastly, she doesn't want us to tell who the next scriber / blogger is because she wants everyone to check the blog. But shhh....guys we are going to help each other out?  (Sorry Ms.Friedmann! I am the one that doesn't obey.)

Lastly, these are the lab results:

The lab results

Lab results with labels

Label each square as either PPT or NR.

PPT= Precipitated

NR=No Reaction

Dissociation and Solvation

Scribe: Juliette Ovadia

Tuesday, October 29th, 2013

Dissociation and Solvation

Agenda

First, we checked in the Dissociation homework from last night and picked up handouts (the power point notes on solutions and ChemThink LabSim Questions). Then we checked the homework and went over questions. Finally, we watched a sugar and salt simulation and went over the notes. 

Dissociation Homework

When going over this homework, Mrs. Friedmann reminded us that dissociations are "breaking apart" reactions. When the reactant dissolves in water, it breaks apart. She reminded us that the products of dissociations are always aqueous because the solute is, in this case, always dissolved in water and breaks apart in water. When making compounds, one should balance the charges, but when separating compounds, one should balance equations. 

For example, in problem #6, there was one molecule, a crystal, of Rb₂SO₄, which had two Rb atoms and one SO_4. This needs to also be shown in the products, so you add a coefficient of 2 in front of the Rb+. A trick is to take the subscript and turn it into a coefficient. We can imagine the Rb₂SO₄  as a molecule breaking apart into different pieces even as we understand that ionic compounds do not actually form these types of molecules.

So, Rb₂SO_{4 (s)}  à 2 Rb⁺ _(aq)  + SO₄^-2_(aq)

Simulation

We wanted to understand what was going on when we dissolved a solute in water, which this simulation (an atomic level picture) demonstrated to us.  The simulation showed a salt shaker filled with NaCl, and we “shook” the salt into the water. We all know that if we shake NaCL in water it dissolves (NaCl is a crystal of an Na ion and a Cl ion, but when mixed with water the Na ions and Cl ions separate). We also know that of things that mix with water, some break apart, and some don’t.

The simulation showed lots of H₂O molecules, and on the molecules were +s and –s and a greek lower case delta. The delta indicates “partial” charge, and the presence of the delta indicates that O₂ has a partial – charge and H₂ has a partial + charge.

The simulation then showed us that H₂O is a polar molecule, meaning hydrogen is positive and oxygen is negative. The simulation showed that when we dissolve NaCl in water, the +s from the hydrogens wrap around the Cl- and the –s from the oxygens wrap around the Na+. Water surrounds the Na+ and Cl- and pulls them apart and dissociates them, because the ability of the H₂O to break the NaCl apart is stronger than the ability of the NaCl to stick together. Other ionic compounds haves +s and –s that have the ability to stick together that is stronger than the ability of H₂O to break them apart. H₂O with salt conducts electricity because the +'s and –'s solvated and conduct electricity.

The process outlined above is called Solvation.

The second part of the simulation showed sugar dissolved in water. Sugar molecules shake apart from each other when dissolved in water but do not dissociate because they do not break into the small pieces that make them up; rather,  only their molecules break apart. Water solvates molecules but not the molecular bonds because the molecular bonds are stronger than the ability of water to break them apart, which is why sugar in water does not conduct electricity.

Answers to questions during the simulation:

•  Each granule of sugar is a huge clump of sugar molecules. Moisture results in the crystals sticking together and crystallizing into a lump of sugar, for instance.
• If one evaporates off the water, a crust of salt would remain because dissociation is physical.
• You can only tell if the ionic compound is stronger than water’s ability to break it apart by testing, and there is a set of rules about what dissolves called the “Solubility Rules” that one must memorize.
• Only ionic compounds are strong electrolytes. Electrolyte: a compound that conducts electricity
• Electrolytes are potassium sodium, calcium and magnesium, and all are ions that perform important jobs in the body. For example, the action potential is when potassium is pushed to one side and chlorine to another to set up a gradient. Electrolytes are necessary to set up gradients in the body!
• Dissolve: to mix at the molecular level (sugar and water dissolve but do not dissociate, while salt dissolves and dissociates.)
•   Remember: If dissociates, must dissolve.

Example: Iced Tea

You cannot sweeten iced tea by simply pouring in sugar because the sugar just settles to the bottom. However, if you heat up the tea, the sugar will dissolve and the tea will taste sweet. Then you would cool the tea back to its former cold temperature. The sugar in the cold iced tea is an example of Slight Solubility, and if you heat a slightly soluble solution like iced tea, the heat will get the sugar to completely dissolve.

Slight Solubility: when you shake something into water and it partially dissolves but some does not dissolve, which was also shown in our simulation.

The sugar in iced tea example uses Colligative Properties that explain the way adding a solute changes the properties of the solvent itself.

Sugar in the iced tea that is heated and dissolves and then cooled results in a Supersaturated Solution, which is delicate and can come out of solution very quickly.

This video shows sodium acetate ("Hot Ice"), a supersaturated solution, crystallizing. 

Example: Rock candy is made with supersaturated sugar water

Homework

ChemThink Lab Simulation

The assignment is on the right side, under the heading Labs, under the heading Chemical Reactions, under the heading Precipitate Lab.

Complete the lab and the ChemThink LabSim worksheet that was handed out.

The next scribe post author is Suvd D. 

Introduction to Unit 4- Solutions & Dissociations

Scribe: Ekene Nwosisi

October 28th 2013

Intro to Solutions & Dissociations 

Agenda

• Received 5 handouts 
• Turned in Micro Mole Rockets Question Sheet that was for HW over the weekend
• Went over Unit 3 Test
• Mini-Experiment (Intro to Electrolytes): 
• Hypothesis: Do all water solutions actually conduct electricity?
• Unit 4 Big Ideas Notes
• Dissociations

Handouts

We received 5 handouts today, which are currently not on Moodle right now (6:00 pm), however, we only got to start one worksheet titled Dissociations (this topic will be further explained below).

Next, we turned in the Micro Mole Rockets Lab Sheet. If you were absent, make sure you turn it in tomorrow.

Unit 3 Test 

• Unfortunately, our class test score average on this test weren't the greatest, as they required deeper thinking and application. However, remember what Ms. Friedmann said last week, it's not the grade that's important, it's how capable we are to handle a challenge that comes at us and know what it takes to solve that problem! 
• Ms. Friedmann then went over the test questions that were confusing to the class

Mini-Experiment 

"What happens when you're at the pool and you see lightning?" Ms. Friedmann asked us. 

The obvious answer to us was to get out because water conducts electricity.

After this, Ms. Friedmann got out a conductivity tester (picture below). She then filled a beaker with water and placed it under the conductivity tester hoping that the light bulb would turn on so that it would prove our hypothesis.

The water was placed under the tester, and to our surprise, the light bulb stayed off. 

Next, Mrs. Friedmann filled another beaker with water that was thoroughly mixed with salt. Once placed under the conductivity tester, the light bulb ignited!

She then tested a beaker filled with water mixed with sugar. Because we were all in the mindset that solutions conducted electricity, we were surprised that the sugar water did not light up the light bulb.

Table Salt (NaCl) mixed
 water conducts electricity

• Explanation of Results
• Can put things in water that can make it conduct electricity (ex: salt)
• The reason that salt, and not sugar, can conduct electricity once mixed with water, is because salt is broken apart in the mixture. The broken apart salt pieces conduct electricity when mixed with water. Sugar, while it can mix with water, does not break apart, and therefore cannot conduct electricity. 
• To further our explanation on this topic, we took notes in our notebooks labeled  "Unit 4 Big Ideas"

Unit 4 Big Ideas Notes:

1. Some things mix with water and some don't
2. Some of the things that mix with water break apart and some don't
3. Breaking apart is called DISSOCIATION 
4. Most things that break apart make it so water can conduct electricity
5. PRECIPITATION reactions happen because something new is made that does not break apart
6. *There are lots of ways to measure and calculate how much stuff is in the water doing the reaction             *(This is the torture part of this unit)

Dissociations

We learned that...

• Dissociations are "breaking apart" reactions
• Physical, not a chemical process
• Only involves ionic compounds

I don't believe that the Dissociations handout we received in class is on Moodle yet, and to those of you that were absent, it is for HW tonight, so the best I can do is put the questions on here, and you can copy them onto your lab notebooks:

Important Notes:

• Break apart ionic compounds as shown in examples 1, 2, 5
• Put the charges of each cation and anion
• Balance equation (look at #5 as an example)

1. NaCl(s)→ Na+(aq)   +   Cl-(aq)
2. KNO3 (s) → K+(aq)     +   NO3- (aq)
3. MgSO4(s)→
4. AgNO3(s)→
5. FeCl3(s)→  Fe3+ (aq)    +  Cl3  (is incorrect b/c no such thing as Cl3) 3Cl1-(aq)
6. Rb2SO4(s)→
7. Zn(CH3COO)2(s)→
8. (NH4)2CO3(s)→
9. Al(NO3)3(s)→
10. MgCrO4(s)→
11. Dissolve potassium hydrocarbonate
12. Dissolve barium acetate
13. Dissolve copper (II) sulfate
14. Dissolve lithium phosphate
    15.  → Pb2+(aq)    +   NO31-

Homework

Complete Dissociations worksheet in lab notebooks

Next Scribe:

Juliette Ovadia