December 4, 2009 (molecular formulas)

Finding Molecular Formulas
It is easy to do if you know the empirical formula and the Molar Mass.

If the empirical formula for a compound is CH2O and the molar mass is 90.0g/mol. Find the molecular formula.
Stock Solutions

• Are concentrated solutions from suppliers

12 M HCl
15 M HNO3
18M H2SO4

Makiki wants to make 200mL of 0.40M HCl. What volume of 10.0M HCl should she use?
V1= ? V2= 0.200L C2V2 = C1V1
C1= 10.0M C2= 0.40M (0.40M)(0.200L) = (10.0M) V1
V1 = 0.008L/ 8mL

Giving Directions - Dec2 Class

In today's class, we learned about which procedures we should use when it comes to concentrations. We also did a pre-lab about how much water we would need to dissolve Copper Chloride completely. We will be doing the lab this friday, Dec4.

Outline for experimental procedures:
- Find the mass you need

Johnny is asked to make 0.95M solution of K2S04. If he needs 300mL, what procedure should he use?
Con'c -> moles -> mass

Give directions to make 3.5 L of 4.5M NaOH

DILUTIONS OF SOLUTIONS
- When you add water, con'c decreases. If the volume is doubled, con'c is halved.

N1 = N2
C1V1=C2V2
Lauren adds 215.0 mL of water to 25.0 mL of 0.8M HCl. Find [HCl]
Janice adds water to 75.0 mL of 0.15M HF to a final volume of 310mL. Find the [HF]
Nicola dilutes 75.0 mL of 0.30M HNO3 to 0.10M. What is the final volume?

How much water did she add?
225mL-75mL = 150mL

Concentration- Nov. 30

In Tuesday's class we learned about concentration and its units.

Vocabulary:
Solution- A homogeneous mixture
Solute- The one present in smaller amount
Solvent- The one present in greater amount
Concentration- Amount of solute/Amount of solvent

Some units for Concentration:
g/mL, g/L, mg/L, mg/mL, ug/L

The most common (and useful) units are:
mol/L = Molarity = Concentration

Examples:
Stefano dissolves 108g of NaBr in enough water to make 300mL of the solution. What is the concentration?
Con'c = 108g/300 mL = 0.36g/mol [NaBr] = ? 0.300L = V

108g x 1 mol/108g M = 1.0mol/0.300L = 3.33mol/ L = 3.33M

Jayson wants to make 600.0mL of 0.60M CaCl2. What mass of solid CaCl2 is required?
V = 0.600L con'c ---> mol ---> mass
mol = M x L = 0.6m/L x 600L = 0.36mol x 111.1g/1mol = 39.996 or 40g

November 16th Class

On November 16th we learned about empirical formula's and how to find them.

Empirical Formula's
- give the whole number of ratio's of elements in a compound while molecular formula's give the actual number.

After learning what empirical formula's are we learned how to find them.

Finding Empirical Formula's

Example:
A sample of an unknown compound is analyzed and found to contain 8.4 g of carbon 2.1 g of hydrogen and 5.6 g of oxygen.

* The easiest way to answer this question is to organize the information that is given into a chart format.


C2H6O

If the ratio ends in a decimal of... 0.5 then multiply by 2.
If the ratio ends in a decimal of... 0.33 or 0.66 then multiply by 3.
If the ratio ends in a deciaml of... 0.25 or 0.75 then multiply by 4.
If the ratio ends in a decimal of... 0.2, 0.4 or 0.6 then multiply by 5.

Nov.12 Class

In the beginning of class we reviewed 10-2 questions and everything up to the mole conversion is on the midterm.

molar mass -> g/mol

molar volume -> L/mol -> mL/mol

Example:
A sample of an unknown gas contains 0.0554mol and occupies a volume of 602.0mL. Determine the molar volume.

Percent Mass of Elements in Compunds

Find the % carbon by mass in Glucose. (C6H12O6)
C = 6(12.o) + H = 12(1.0) + O = 6(16.0) = 180 g/mol
% of Carbon = 72g/180g = 0.4 -> 40%
% of Hydrogen = 12g/180g = 0.0666... -> 7%
% of Oxygen = 96g/180g = 0.5333... -> 53%

Percent Composition
- means the % mass of each element in a compound.

Find the composition of NO3
N = 1(14.0) + O = 3(16.0) = 62.0g/mol
% of Nitrogen = 14g/62g = 0.228 -> 23%
% of Oxygen = 48g/62g = 0.77 -> 77%

Finding the mass of an element in a given mass of a compound.

Find the mass of Carbon contained in a 30.0g sample of CO3.
(First find the percent. Then take the percentage as a decimal and multiply it by the given amount.)
C = 1(12.0) + O = 3(16.0) = 60g/mol
% C = 12g/60g = (0.200) (30.0) = 6g
% O = 48g/60g = (0.800)(30.0) = 24g

Find the mass of K,C, and O contained in K2CO3 if the sample is 500.0g
K = 2(39.1) = 78.2 + C = 12 + O = 16(3) = 48 = 138.2g/mol
% of K = 78.2/138.2 = 56.6% -> 0.566(500g) = 283.0g
% of C = 12/ 138.2 = 8.7% -> 0.087(500g) = 43.5g
% of O = 48/138.2 = 34.7% -> 0.347(500g) = 173.5g
Total = 500.0g

Density & Moles - Continued: Nov9 Class

Today in class, we drew out the flow chart for all the calculations we could do with this unit of moles. We did 3 examples and the last one is one made up.

Example: 1.25L of an unknown gas has a mass of 3.47g. What is the molar mass if it is at STP?
Example 2: A 5.00g sample of solid lead contains 0.247 mol of Pb. Calculate the density of Lead.

Example 3: 250mL of a gas which is known to contain one sulfur atom and an unknown number of fluorides has a mass of 1.63 g at STP
SFx
Step 1: Find the molar mass (g/mol)

Step 2: Find the chemical forumla


Example 4: A 8.00mL sample of (solid) Tin contains 0.632 mol of Sn. Calculate the density of Tin.


We spent the last 20-25 minutes of class, finishing questions #11-30 on the worksheet.

Density & Moles- Nov 5

On Thursday class we learned about density of gases at STP and how to convert into mole from different units.

Density ---> mass per unit volume


Density of gases at STP
-1 mole of gas

Calculate the density of O2 at STP


Example:

A mystery gas has a density of 1.696g/L at STP. It is a diatomic element. Identify the gas.
MM = (1.696 g/L)(22.4 L/mol) = 38 g/mol--->19 g/mol---> Flourine, F2


The density of Boron (solid) is 2.34 g/mL. How many molecules are in a 60.0 mL piece?



Liquid Mercury has a density of 13.55 g/mol. Find the volume occupied by 1.806 x 10^27 atoms of G

Atoms & Molecules- Oct. 30

On Friday we learned about Atoms & Molecules:



For mono atomic elements
a molecule = an element
Eg. Ne=Ne

Diatomic elements
Molecule, Element
Cl2, Cl

Molecules of compounds
H-O-H


Molecule
2 'H' atoms in 1 molecule
1 'O' atoms in 1 molecule


Example:
Write Ammonium Carbonate ---> (NH4)2 CO3
N: 2
H: 8
O: 3
C: 1

Moles <-----> Molecules
6.02 x 10^23 molec/ 1 mol


Example:
How many molecules are in 0.25 mol of CO2?
0.25 mol x 6.02 x 10^23/ 1 mol = 1.505 x 10^23 molec


Example:
5.1772 x 10^24 molecules of H2O = ? moles
5.1772 x 10^24 molecules 1 mol/6.02 x 10^23 =


Example:
Find the number of 'H' atoms in 4.0 mol of ammonia(NH3)?
moles--->molecules--->H atoms
4.0 mol x 6.02 x 10^23/ 1 mol = 2.41 x 10^24 molecules x 3 = 7.22 x 10^24 'H' atoms

Mole Ratio Lab - October 28 Class

Today in class, we completed the Mole Ratio Lab. Our objective of the lab was to find out the ratio of moles of iron, to moles of copper.

The Pre-Lab consisted of:
1) 2Fe + 3CuCl2 -> 3Cu 2FeCl3.
2) The ratio of moles of copper produced to moles of iron consumed is 3:2

The materials we used were:
Apparatus:
Beakers (250 mL)
Wash bottle
Stirring Rod
Crucible tongs
Centigram balance
drying oven
safety glasses
lab apron
plastic gloves
sand paper or emery cloth
face shield
Reagents:
copper (II) chloride
2 iron nails (approx. 5 cm)
1 M hydrochloric acid
distilled water

Our procedure:
1. Find the mass of 250mL beaker. Record the mass to the nearest hundredth of a decimal (g)
2. Add 8g of copper (II) chloride crystals to the beaker. Find and record the mass.
3. Add 50mL of distilled water to the beaker. To dissolve the crystals, swirl the beaker around.
4. Clean and dry 2 nails. Use sand paper if needed. Then, find and record the mass of the nails.
5. Put the nails into the solution and leave them there for 20 min. Observe the formation of copper and some of the iron that will be used up in the beaker.
6. Pick up the nails one by one using the tongs. Before removing the nails from the beaker, use distilled water to rinse off any remaining copper. Use a stirring rod to scrape any excess copper if needed. Let the nails dry on a paper towel.
7. Find and record the mass of the nails after they are completely dry.
8. (Decant- pour off only the liquid from a container that is holding both liquid and solid.)
Decant the liquid from the solid by putting the liquid into another beaker.
9. After decanting, rinse the solid with 25mL of distilled water. Decant again. Repeat this step 4 more times.
10. Wash the solid with 25mL of 1M hydrochloric acid. Decant again (twice). Then clean the solid with 25mL of distilled water.
11. Place the oven in a drying oven to dry.
12. Let the copper dry, then find and record the mass of the beaker + copper.
13. Make sure to clean up your lab and wash your hands properly.

Our observations:
Mass of empty dry beaker = 159.34g
Mass of beaker + copper (II) chloride = 167.34g
Mass of 2 iron nails = 5.45g
Mass of 2 iron nails (after) = 4.95g / 5.0g
Mass of beaker with copper (after) = 160.72g

We discovered that 0.5g of iron was used in the reaction and 4.95 g of copper was produced.
Our final ratio was 1.5:1 and we had a percent error of 56%. This is because we could have spilled some of the liquids and there was too much liquid in the copper and it didn't dry out fully.

Gases & Moles - Oct. 26 Class

Gases and Moles

The volume of a balloon occupied by a certain gas depends on the temperature and pressure.

Standard Pressure & Temperature (STP)

* 0°C & 101.3 kPa

(273 K)

Standard Ambient Temperature & Pressure (SATP) (24.8L/mol)
* 25°C & 100kPa

(298 K)

The volume of 1.0 mole of any gas at STP is 22.4

The molar volume at STP is 22.4L

Example

1) Find the volume occupied by 0.060 mol of CO2 gas at STP

2) Find the number of moles in a 264.0 mL sample of NO2 at 0° and 101.3 kPa(STP)

3) Find the volume occupied by 22.0g of CO2 (g) at STP

Atomic Mass - October 21st Class

Today, in class we learned how to find the molar mass in compounds.
Atomic Mass: The mass of 1 mole of atoms in an element
- The mass of 1.0 mol of 'C' atoms is 12.0g
- The mass of 1.0 mol of 'Ca' atoms is 40.1g

Molecular Mass: The mass of 1.0 mole of molecules of an element or compound
N2, O2, F2, Br2, H2, Cl2, I2
P4, S8

Assume all the rest are monoatomic



Finding the Molar Mass of Compounds
H2O
2 H= 2.0 (1.0) = 2.0
1 O = 1 (16.0) = 16.0
Total = 18.0 g/mol



- Find the molar mass of Ammonium phosphate

NH4+
PO4³- = (NH4)3PO4

3 N = 3(14.0) = 42

12 H = 12(1.0) = 12

1 P = 1(31.0) = 31

4 0 = 4(16.0) = 64

Total = 149 g/mol



Converting Mass <-> Moles

The Mole - Oct. 19 Class

In the beginning of class, Mr. Doktor showed us this really cool experiment involving this equation: 2 H2 + O2 -> 2 H2O

This is known as the Hydrogen bomb equation.





To count 1 mole, it will take:
6.03 x 10^23 = 1.0 x 10^22 mins
= 1.67 x 10^20 hours
= 6.97 x 10^18 days
= 2.3 x 10^17 months
= 1.94 x 10^16 years

The mole

1 mole = 602 000 000 000 000 000 000 000 = 6.02 x 10^23 -> Avogradro's number

2 H2 + O2 -> 2H20

2 H2 molecules + 1 O2 molecule -> 2 molecules of H2O

12.04 x 10^23 + 6.02 x 10^23 -> 12.04 x 10^23

of H2 molecules molecules of O2 molecules of water

How big is Avogadro's Number?

$1 mol

$6.02 x 10^23

6.0 x 10^9 = population of earth

$ 6.0 x 10^23

__________

6.0 x 10^9 ppl

= $1.0 x 10^14 -> $100 000 000 000 000

How gases combine

John Dalton

-look at masses of gases

11.1g of H2 reacts with 22.9g of O2

46.7g of N2 reacts with 53.3g of O2

42.9g of C reacts with 57.1g of O2

= No pattern

Joseph Gay Lussac

-combine gases based on volume

1L of H2 reacts with 1L of Cl2 -> 2L of HCl H2 + Cl -> 2 HCl

1L of N2 reacts with 3L of H2 -> 2L NH3

2L of CO reacts with 1L of O2 -> 2L CO2

= gases combine in simple whole number ratios

Avogadro's Hypothesis

-Equal volumes of any gas at a constant temperature and pressure contain equal numbers of molecules.

H2

same / but different mass

02

In conclusion, at the end of class, we did another really cool experiment known as the potato gun! It needs oxygen to react with chemicals.

Hydrate Lab

Today in class we learned tht hydrates are ionic compounds that contain an inorganic salt compound loosely bound to water. Putting our knowledge into action we did an experiment today that was meant to determine the emprical formula of a hydrate. In the lab we determined the anhydrous (without water) mass of the hydrate and then compared the original mass with the actual mass of water that should be present.

The materials that we used to conduct this experiment are:

- bunsen burner

- test tubes

- test tube rack

- test tube clamp

- weight scales

The first step we took in conducting our experiment is we filled a test tube with about 1 cm of the hydrate. We then carefully placed the test tube on the scale and recorded the mass of the hydrate and test tube. With extreme cation we proceeded to connect and light our Bunsen burner and adjusted the gas flow until the flame was about 5 cm tall. Afte heating our test tube over the Bunsen burner with the clamps (in and out) for about 5 minutes we carefully re-weighed the test tube.

Our observations were that the mass before heating was:

and the mass after heating was:

Our conclusion was that # amount of water was released during heating and that # percent of the hydrate was water. The actual percent ofwater in the hydrate was 45%. Our percent error was #.

Overall we learned a lot from the first-hand experience and had a great time!

Acids & Bases- Oct. 8 Class

In today's class we learned the characteristics of acids and bases.

Acids
- Solid, Liquid, or gas at SATP (25 degrees celcius, 100 kPa)
- Form conducting aqueous solutions
- Turn blue litmus red
- Dissolve in water to produce H+
- Taste sour


Bases
- Turn red litmus blue
- Slippery
- Non-conductive
- Dissolve in water to produce OH-




Naming Acids
- Acids are aqueous (dissolve in water)
- Hydrogen compounds are acids

• HCL(aq) --> Hydrochloric
• AcidH2SO4(aq) --> Sulfuric Acid
  

- Hydrogen appears first in the formula unless it is part of a polyatomic group

• CH3COOH(aq) --> Acetic Acid

- Classical rules use the suffix "ic" and/or the prefix "hydro"

• Eg. Hydrochloric Acid

- IUPAC system uses the aqueous hydrogen compound

• Eg. HCl(aq) --> Aqueous Hydrogen Chloride

Naming Bases
- For now, all bases will be aqueous solutions of ionic hydroxides

• NaOH
• Ba(OH)2
  

- Use the cation name followed by hydroxide

• Sodium Hydroxide
  
• Barium Hydroxide
  

Examples
- HI(aq) --> Hydrochloric Acid
- H3PO4(aq) --> Phosphoric Acid
- H3PO4(aq) --> Phosphorous Acid
- HNO3(aq) --> Nitric Acid
- HNO2(aq) --> Nitrous Acid
- Mg(OH)2 --> Magnesium Hyrdroxide
- HBr(aq) --> Hydrobromic Acid
- HOOCCOOH(aq) -- Oxalic Acid


How to make your own PH indicator between an acid and base!

October 6th 2009

Today in chemistry class Mr. Doktor showed us ...

Hydrates

- Some compounds can form lattices that bond to water molecules
Example: Copper Sulfate
- These crystals contain water inside them which can be released by heating
- Without water the compound is often preceded by “anhydrous”

Naming Hydrates

1. Write the name of the chemical formula.
2. Add a prefix indicating the number of water molecules.
3. Add hydrate after the prefix

Molecular Compound

- Composed of two or more non-metals.
- Low melting point and boiling point.
- Share (not exchange) electrons
- Usually end in –gen or –ine
Example: Hydrogen, Oxygen…
- 7 molecules are diatomic
They are… Hydrogen, Nitrogen, Oxygen, Fluorine, Cholrine, Bromine and Iodine
Two molecules are polyatomics.

Chemical Nomenclature - Oct 2 Class

In the beginning of class, Mr. Doktor showed us an example of electrolysis. He used pickles, water, and a battery charger. He then touched the pickle with the two wires, and you could see smoke come out and the pickle being electricuted. He then touched the water with the wires and smoke came out as well. You could see the hydrogen and oxygen as the smoke and you could smell this funky burning smell. You should not try this experiment at home due to dangerous results such as being electricuted or starting a fire.



Here is a video of an example of electrolysis.

http://www.youtube.com/watch?v=Or22ktW8btc&feature=related
This is a quick video that shows how hydrogen and oxygen atoms blow up a balloon. In my assumption, i believe that the balloon blew up due to evaporating water.



After that experiment in class, Mr. Doktor gave us examples and questions about naming compounds. Here are the examples done in class:
Name:
1. PbS2 - Lead (IV) Sulphide
2. MgO - Magnesium Oxide
3. CuCl2 - Copper (II) Chloride
4. Cr2O3 - Chromium (III) Oxide

Chemical Nomenclature
- Naming chemical compounds has been a very difficult task and different systems have been used through the centuries.
- Today the most common system is IUPAC for most chemicals.
- Ions
* Binary Ionic
- Polyatomic Ions
- Molecular Compounds
- Acids



Chemical Formulas
Be aware of the differences between ion and compound formulas.
Zn²+ - ion charge
BaCl2 - number of ions (Subscript)

Naming Ions
- For metals, use the name of the element and add ion
Eg. Al³+ = Aluminium Ion
- For non-metals, remove the original ending and add -ide
Eg. F- = Fluorine -> Fluoride

Polyatomic Ions have special names.

Binary Ionic
Steps:
1. Write the formula for the cation first and then the formula for the anion. (Cation = positive because "cat" = "pussytive" & Anion = Negative)
2. Criss cross charges moving the numbers below.
3. Reduce ion numbers to lowest common multiples, omit 1 and omit charges.
Eg.


Example
- Write the chemical formulas of:
Aluminum Fluoride = AlF3
Sodium Oxide = Na2O
Iron (III) Sulphide = Fe2S3

Multivalent Ions
- Some elements can form more than one Ion
Eg. Iron -> Fe³+ or Fe²+ / Copper -> Cu²+ or Cu¹+
- The more common ion is the top one of the P.T.
- IUPAC uses roman numerals in the parenthesis to show the charge.
- Classical (ie old) systems uses latin names of elements and the suffixes -ic (larger charge) and -ous (smaller charge)
FeO -> Ferrous Oxide
Fe2O3 -> Ferric Oxide

Other Classical Names
- Ferr = Iron
- Cupp = Copper
- Mercur = Mercury
- Stann = Tin
- Aunn = Gold
- Plumb = Lead
Chemical Formula: Sodium Nitrate -> NaNO3 (Na+/NO3-)
Barium Phosphate -> Ba3(PO4)2

Classification of Matter - September 30 Class

In this class, we learned about Homogeneous and Heterogeneous substances.. We distinguished that tap water is a heterogeneous substance because there are other chemicals that are put into the tap water when running through the drains, sinks, etc. Mr. Doktor also passed around a beaker filled with aluminum pieces and salt, a heterogeneous mixture that didn't react together. We also learned that there are certain ways of separating mixtures, which are mostly physical changes.

Notes:
Classification of Matter

• Understanding matter begins with how we name it. We can divide matter into two parts: Homogeneous substances and heterogeneous substances
  
• Homogeneous: consists of only one visible component
  - distilled water, oxygen, graphite
  
  
  
• Heterogeneous: contain more than one visible component
  - chocolate chip cookie, granite

Pure Substances:

• There are 2 types of Pure Substances:
  Element: substances that cannot be broken down into simple substances by chemical reactions
  - oxygen, iron, magnesium
• Compound: substances thar are made up of 2 or more elements & can change into elements (or other compounds by chemical reactions)
  - water, sugar

Telling the difference

• It is often very difficult to know if something is an element or a compound
  - The differences are only "visible" on the atomic level
• One method is to connect the substance to an electric current. This technique. called electrolysis can split the compound apart into its constituent compounds

Solution

• A solution is a homogeneous mixture of 2 or more substances
  - Solutions usually involve but don't have to (fog, steel)
• The compound present in greater mount is the solvent
  - Water is the most ommon solvent
  - The symbol (aq) is used when something is dissolved in water
• The compound present in smaller amount is the solute
  - In salt water, salt is the solvent

Mixtures

• Many mixtures are easy to identify (chocolate chip cookies) but others are easily confused as pure substances
• In heterogeneous mixtures.. the different parts are clearly visible (granite, sand)
• In homogenous mixtures.. the different parts are NOT visible (salt, water, air, brass)

Separating Mixtures

• There are many methods to separate mixtures, depending on how the type of mixture
  - By Hand (Heterogenous mixture only)
  - Filtration (Heterogeneous mixture only)
  - Distillation
  - Crystallization
  - Chromatography
• All of these are physical changes

Matter and its Changes- Sept 28 Class

In Monday's lesson we learned the States of Matter, Changes in Matter, Physical Changes, Phase Changes, Chemical Changes, and the Conservation of Matter. Also, we got to watch a youtube video of mythbusters as an example to figure out the chemical and physical changes that occurred.


What is Matter?
-Anything that has mass and occupies space
-Matter can exist in many different states, the most common are:

• Solid, Liquid, Gas
  
• Plasma, Aqueous (dissolves in water), Amorphous
  
  

-Solids: holds one shape and has a definite volume (strong bonds)
-Liquids: can change shape, but has a definite volume (weak bonds)
-Gas: can change shape & volume (no bonds)
-Aqueous: something dissolved in water
-Plasma --> ionization


Changes in Matter
- Matter can undergo many changes
- Nearly all changes can be broken down into 3 categories:

• Physical Changes
  
• Chemical Changes
  
• Nuclear Changes
  

Physical Changes
- Involves changing shape or state of matter (crushing, tearing, etc.)
- No new substances are formed (Eg. Boiling water, cutting wood,
smashing cars)




Phase Changes
- Changing from a solid to a gas can often be confused as a chemical change
- Chemicals remain the same


Chemical Changes
- New substances are formed
- Properties of the matter change

• Conductivity, acidity, color, etc.
• Eg. Iron rusting, burning wood, digested food

Physical change of water into hydrogen peroxide


Conservation of Matter
- In Physical and Chemical Changes, matter is neither created nor destroyed.

This is the link of the video we watched of mythbusters:
http://www.youtube.com/watch?v=MWJU6sbf8Ng

Chapter 1 Review- September 21st Class

Today in Class we finished doing our NaCl labs and write-ups. We also had time to review for the upcoming chapter one test.

Below are some helpful study tools to study the various things in chapter one that we will be tested on.

Sodium Chloride Lab - Sept 17 Class

In the beginning of class, we reviewed Dimensional Analysis since we didn't get a chance to go over it thoroughly in the previous class. Here are some examples of unit conversions:


If you have trouble converting milligrams into kilograms, then you could always break it down into grams first. The 2 kilograms and the other 2 grams cancel each other out, leaving you with 120.000,000mg (120,000 x 1000).


This problem includes dealing with measurements and time as well. We must remember to incorporate both information. The Litres and the seconds each have another pair to cancel each other out, leaving you with 435 000mL/min (7.25 x 60 x 1000).


Once again we have broken down kg into g because it is easier to convert g into Mg. This problem also includes time that needs conversion as well. The kilograms, seconds, and grams all cancel out, and the end result is 626.4Mg/h (174 x 3600 x 1000/1,000,000).


After we did a little bit of review, we started our very first lab called Sodium Chloride.


















Purpose
To experiment the maximum amount of salt you can dissolve in 10mL, 20mL, 30mL, and 40mL.

Materials
- Sodium Chloride
- 150mL of Distilled Water
- 100mL graduate cylinder
- 50mL beaker
- 100mL beaker
- Weight paper
- Electronic Scale
- Lab coat
- Safety Glasses

Procedure
1. Gather all the materials an put them on the lab bench.
2. Measure 10mL of distilled water using the graduated cylinder. Transfer this water into a 50mL beaker.
3. Weigh 50 g of Sodium Chloride.
4. Add Sodium Chloride to the water until it stop dissolving and the first salt crystals begin appearing on the bottom of the beaker.
5. Measure the mass of salt remaining. Record the difference in salt as the amount added to the 10mL of water.
6. Repeat steps 2-5 for 20mL, 30mL, and 40mL of water.
7. Record all your data in the table below.
8. Create a graph of Mass of Salt vs Volume of Water.

Observations

The results can and may vary.





http://www.youtube.com/watch?v=Kxbc8nuv_0k

This video is similar to the lab we did today. The only differences were the scales, how they measured the ingredients, and they shook the salt and water together instead of stirring it with a glass rod.