In this class, we did a lab to experimentally determine the molar volume of a gas.We had to measure how much mass does butane have and also the molar mass.
The directions for the lab is as follows:
1. Set up the weigh scale.
2. Weigh the mass of the Butane lighter.
3. Fill the sink with water. Make sure to clog it so that water won’t drain.
4. Fill the graduated cylinder with water.
5. Turn it upside down while keeping the water trapped inside.
6. Direct the Lighter under the graduated cylinder.
7. Fill the graduated cylinder with about 20-25mL of Butane.
8. Dry the Butane lighter after use.
9. Weigh the mass of the Butane lighter.
10. Calculate the mass of Butane that was used.
11. Calculate how many moles of Butane are in the mass. ( )
12. Determine the Molar Volume of Butane. ( )
Tuesday 29 November 2011
Multistep Conversions!
Multistep conversions can be used to figure out the mass, moles, volume and even molecules and atoms in a certain chemical equation.
Mass to Volume:
Step 1: Convert mass to moles.
Step 2: Convert moles to volume.
Mass to Molecules:
Step 1: Convert mass to moles.
Step 2: Convert moles to molecules.
Mass to Atoms:
Step 1: Convert mass to moles.
Step 2: Convert moles to molecules.
Step 3: Convert molecules to atoms.
Volume to Mass:
Step 1: Convert volume to moles.
Step 2: Convert moles to mass.
Volume to Molecules:
Step 1: Convert volume to moles.
Step 2: Convert moles to molecules.
Volume to Atoms:
Step 1: Convert volume to moles.
Step 2: Convert moles to molecules.
Step 3: Convert molecules to atoms.
Examples:
1. 5.1g x 1mol/79.9g x (6.02x10^23)/1mol = 3.8 x 10^22
2. 3.62x10^24 x 1mol/(6.02x10^23) x 320g/1mol = 192g
3. 2.94x10^24 x 1mol/(6.02x10^23) x 142g/1mol = 693g
Mass to Volume:
Step 1: Convert mass to moles.
Step 2: Convert moles to volume.
Mass to Molecules:
Step 1: Convert mass to moles.
Step 2: Convert moles to molecules.
Mass to Atoms:
Step 1: Convert mass to moles.
Step 2: Convert moles to molecules.
Step 3: Convert molecules to atoms.
Volume to Mass:
Step 1: Convert volume to moles.
Step 2: Convert moles to mass.
Volume to Molecules:
Step 1: Convert volume to moles.
Step 2: Convert moles to molecules.
Volume to Atoms:
Step 1: Convert volume to moles.
Step 2: Convert moles to molecules.
Step 3: Convert molecules to atoms.
Examples:
1. 5.1g x 1mol/79.9g x (6.02x10^23)/1mol = 3.8 x 10^22
2. 3.62x10^24 x 1mol/(6.02x10^23) x 320g/1mol = 192g
3. 2.94x10^24 x 1mol/(6.02x10^23) x 142g/1mol = 693g
Sunday 27 November 2011
Converting between Atoms/Molecules & Moles!
It takes one step to convert from Moles to Molecules and two steps from Moles to Atoms.
Moles to Molecules:
The conversion factor for moles to molecules is Avogadro's number (6.02 x 10^2)
Example:
How many water molecules are there in 0.65mol?
Step 1: Lay down conversion equation.
0.65mol x 6.02x10^23molec / 1mol = ???
Step 2: Multiply by Avogadro's number and divide by 1mol (to cancel the mol).
0.65mol x 6.02x10^23molec / 1mol = 3.9x10^23 molec
Moles to Atoms:
Convert Moles to Molecules first by using Avogadro's number as a conversion factor then use subscripts
Example:
How many oxygen atoms are in 3.9x10^23molec?
Step 1: Lay down conversion equation.
3.9x10^23molec x 2atoms/1mol = ???
Step 2: Multiply the molecules by the number of subscripts.
3.9x10^23molec x 2atoms/1mol = 7.8x10^30 oxygen atoms.
Moles to Molecules:
The conversion factor for moles to molecules is Avogadro's number (6.02 x 10^2)
Example:
How many water molecules are there in 0.65mol?
Step 1: Lay down conversion equation.
0.65mol x 6.02x10^23molec / 1mol = ???
Step 2: Multiply by Avogadro's number and divide by 1mol (to cancel the mol).
0.65mol x 6.02x10^23molec / 1mol = 3.9x10^23 molec
Moles to Atoms:
Convert Moles to Molecules first by using Avogadro's number as a conversion factor then use subscripts
Example:
How many oxygen atoms are in 3.9x10^23molec?
Step 1: Lay down conversion equation.
3.9x10^23molec x 2atoms/1mol = ???
Step 2: Multiply the molecules by the number of subscripts.
3.9x10^23molec x 2atoms/1mol = 7.8x10^30 oxygen atoms.
Monday 21 November 2011
Moles to Volume Conversions
The volume of 1 mol in a substance is molar volume.
At a specific temperature and pressure one mole of any gas occupies at the same volume.
- At 0 °C and 101.3kPa 1 mol = 22.4 L
- This temperature and pressure is called STP
* 22.4 L/mol is the molar volume at STP (Standard Temperature Pressure)
Examples:
At a specific temperature and pressure one mole of any gas occupies at the same volume.
- At 0 °C and 101.3kPa 1 mol = 22.4 L
- This temperature and pressure is called STP
* 22.4 L/mol is the molar volume at STP (Standard Temperature Pressure)
Examples:
Converting from Moles to Mass!
Mr. Doktor explained to us how to convert Moles into Mass.
The mole being converted is multiplied by mass/1 mol.
Ex:
0.89mol x 111.1g/1mol = 98.88g
1.112mol x 20.0g/1mol = 22.2g
0.159mol x 60.1g/1mol = 9.56g
Mr Doktor also showed us how to convert Mass into Moles.
The mass being converted is multiplied by 1 unit of mass/mol.
Ex:
158.1g x 1mol/303.3g = 0.5213mol
362.8g x 1mol/72g = 5.04mol
12.35g x 1mol/58.0g = 0.140mol
The 1mol is either on the top or bottom so that units can cancel.
The mole being converted is multiplied by mass/1 mol.
Ex:
0.89mol x 111.1g/1mol = 98.88g
1.112mol x 20.0g/1mol = 22.2g
0.159mol x 60.1g/1mol = 9.56g
Mr Doktor also showed us how to convert Mass into Moles.
The mass being converted is multiplied by 1 unit of mass/mol.
Ex:
158.1g x 1mol/303.3g = 0.5213mol
362.8g x 1mol/72g = 5.04mol
12.35g x 1mol/58.0g = 0.140mol
The 1mol is either on the top or bottom so that units can cancel.
Thursday 17 November 2011
Molar Mass!
Today during chemistry class, Mr. Doktor taught us how to find the molar mass of a substance and a compound.
•The mass (in grams) of 1 mole of a substance if called the molar mass.
•It can be determined from the atomic mass on the periodic table.
•It is measure in g/mol
Ex. Iron(Fe) – 55.8 g/mol
Ex. Oxygen (O2) – 2(16) = 32.0 g/mol
Molar Mass of Compounds
• To determine the molar mass of a compound add the mass of all the atoms together.
Ex. NaCl-------- 23 + 35.5 = 58.5 g/mol
Ex. AlCl3------- 27 + 3(35.5) = 133.5 g/mol
•The mass (in grams) of 1 mole of a substance if called the molar mass.
•It can be determined from the atomic mass on the periodic table.
•It is measure in g/mol
Ex. Iron(Fe) – 55.8 g/mol
Ex. Oxygen (O2) – 2(16) = 32.0 g/mol
Molar Mass of Compounds
• To determine the molar mass of a compound add the mass of all the atoms together.
Ex. NaCl-------- 23 + 35.5 = 58.5 g/mol
Ex. AlCl3------- 27 + 3(35.5) = 133.5 g/mol
Wednesday 9 November 2011
Avagadro's Number (How we count atoms)
Today in chemistry class, Mr Doktor taught us how to calculate how many atoms there in a mole of atoms using the Avagradro's number, we learn that:
-Atoms and molecules are extremely small
-Macroscopic bjects contain too many atoms to count or weigh individually
-Amedeo Avogadro proposed that the number of atoms in 12.00000g of carbon be equal to a constant
-This value is now called Avogadro's number and forms te basis of all quantitative chemistry
-1 mol= 6.02x10^23
-1 pair = 2, 1 dozen = 12, 1 century = 100, 1 mol= 6.02x10^23
e.g. a sample of carbon contains 2.47 x10^25 atoms. How many moles of carbon is this?
-Atoms and molecules are extremely small
-Macroscopic bjects contain too many atoms to count or weigh individually
-Amedeo Avogadro proposed that the number of atoms in 12.00000g of carbon be equal to a constant
-This value is now called Avogadro's number and forms te basis of all quantitative chemistry
-1 mol= 6.02x10^23
-1 pair = 2, 1 dozen = 12, 1 century = 100, 1 mol= 6.02x10^23
e.g. a sample of carbon contains 2.47 x10^25 atoms. How many moles of carbon is this?
Sunday 6 November 2011
Hydrate Lab
Todays class we did a Lab on Hydrates. We were split into groups of two or three to complete the lab. The purpose of our Lab was to measure how much water was in the hydrate.
We were first taught to use a Bunsen Burner, and how to hold the test tube above the bunsen burner so that we wouldn't hurt ourselves. Before we put any hydrate into our test tube, it is to be measured on to the scale. Next, we would receive some hydrate and put in about 1cm into our test tube, put it on the scale and measure how many grams it is all in all. We then would hold up the test tube above the Bunsen Burner until all the water evaporates, it is then put back onto the scale and measured.
After we have gotten all the measurements needed, we would have to find it's percent error using the formula:
To know you have done the Lab correctly, your percent error should've ranged somewhere below 10%
We were first taught to use a Bunsen Burner, and how to hold the test tube above the bunsen burner so that we wouldn't hurt ourselves. Before we put any hydrate into our test tube, it is to be measured on to the scale. Next, we would receive some hydrate and put in about 1cm into our test tube, put it on the scale and measure how many grams it is all in all. We then would hold up the test tube above the Bunsen Burner until all the water evaporates, it is then put back onto the scale and measured.
After we have gotten all the measurements needed, we would have to find it's percent error using the formula:
To know you have done the Lab correctly, your percent error should've ranged somewhere below 10%
Wednesday 2 November 2011
Naming Compounds Part 2!
Today Mr Doktor explained to us how to name molecular compounds, acids and bases. There are various rules for naming these compounds:
Molecular Compounds:
- There are 7 Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2
- There are 2 Polyatomic molecules: S8, P4
Rules for naming a molecular compound:
- Use the name of the first element
- Second element ends in -ide
- 1st atom usually does not have a prefix (Ex. NO -> Nitrogen monoxide)
- Hydrogen doesn't have a prefix (Ex. H2S -> Hydrogen sulfide)
- Some compounds are to be memorized:
Naming Acids:
- Hydrogen compounds are acids (Ex. HCl -> Hydrochloric acid, H2SO4 -> Sulphuric acid)
- Hydrogen appears first in the formula unless it is part of a polyatomic group (Ex. CH3OOOH -> Acetic Acid)
- Classical rules use the suffix -ic and/or the prefix hydro- (Ex. Sulphuric acid, Hydrochloric acid)
- IUPAC system uses the aqueus hydrogen compound (Ex. HCl (aq) -> Aqueous Hydrogen Chloride)
Naming Bases:
- Cation and OH (Ex. NaOH, Ba(OH)2)
- Use the cation name followed by "hydroxide" (Ex. Sodium Hydroxide, Barium Hydroxide)
Some Acids and Bases:
Molecular Compounds:
- There are 7 Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2
- There are 2 Polyatomic molecules: S8, P4
Rules for naming a molecular compound:
- Use the name of the first element
- Second element ends in -ide
- 1st atom usually does not have a prefix (Ex. NO -> Nitrogen monoxide)
- Hydrogen doesn't have a prefix (Ex. H2S -> Hydrogen sulfide)
- Some compounds are to be memorized:
| IUPAC Name | Formula |
|---|---|
| Water | H20 |
| Hydrogen Peroxide | H2O2 |
| Ammonia | NH3 |
| Glucose | C6H12O6 |
| Sucrose | C12H22O11 |
| Methane | CH4 |
| Propane | C3H8 |
| Octane | C8H18 |
| Mathanol | CH3OH |
| Ethanol | C2H5OH |
Naming Acids:
- Hydrogen compounds are acids (Ex. HCl -> Hydrochloric acid, H2SO4 -> Sulphuric acid)
- Hydrogen appears first in the formula unless it is part of a polyatomic group (Ex. CH3OOOH -> Acetic Acid)
- Classical rules use the suffix -ic and/or the prefix hydro- (Ex. Sulphuric acid, Hydrochloric acid)
- IUPAC system uses the aqueus hydrogen compound (Ex. HCl (aq) -> Aqueous Hydrogen Chloride)
Naming Bases:
- Cation and OH (Ex. NaOH, Ba(OH)2)
- Use the cation name followed by "hydroxide" (Ex. Sodium Hydroxide, Barium Hydroxide)
Some Acids and Bases:
| Acid/Base | Compound |
|---|---|
| Hydrochloric Acid | HCl |
| Nitric Acid | HNO3 |
| Suphuric Acid | H2SO4 |
| Phosphoric Acid | H3PO4 |
| Acetic Acid | CH3COOH |
| Ammonia | NH3 |
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