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!