When you find a percent composition of molecules, you calculate the molar mass of an element, which is shown under each element on the periodic table, then calculate each element's percentage of that total.
formula = mass of element/ mass of compound x 100 %
Example:
What is the % composition of H2O ?
First, calculate the total molar mass of H2O
Molar mass of 2 hydrogen is 2.0 g/mole . Molar mass of oxygen is 16.0 g/mole.
So, the total molar mass of H2O is 2.0 + 16.0 = 18.0 g/mole
% H = 2.0 / 18.0 x 100% = 11.1 %
% O = 16.0 / 18.0 x 100% = 88.9 %
Empirical Formula: The ratio of the number of elements in a compound expressed in lowest form Ionic Compounds are usually an example of empirical formulas
Example: NaCl
MgO
Molecular Formula: The actual amount of the number of elements in a specific compound Some covalent compound's molecular formula is not the same as their empirical formula Example: C4H10 (The empirical form would be C2H5)H2O2 (The empirical form would be HO)To convert from and Empirical to Molecular or Molecular to Empirical, we need to find the whole number to multiply it by.
Converting Between Empirical and Molecular Formulas:
The three formulas are:
Molecular Formula = Empirical × Whole Number (N)
Mass of 1 mole = Empirical mass in grams × Whole Number
Molecular Formula Mass = Empirical × Whole Number
If the empirical formula of a compound is C2H5and the molar mass is 58g/mol, what is the molecular formula for this compound? First, we find the mass the C2H5 (Empirical Formula) C = 12.0x2 = 24.0g H = 1.0x 5 = 5.0g 24.0+5.0 = 29.0g
Then, divide the molar mass by the mass of the empirical formula to calculate the whole number (N) used to separate between the emiprical and the molecular formulas.
58g/mol ÷29.0 g = 2
Find the empirical formula of a compound that contains 7.2grams of Carbon, 1.2grams of Hydrogen, and 9.6grams of Oxygen.
Step 1: Convert the grams of elements into moles
C: 7.2g C×1 mol C/12.0g C = 0.6 mol
H: 1.2g H×1 mol H/1.0g H = 1.2 mol
O: 9.6g O×1 mol O/16.0g O = 0.6 mol
Step 2: Divide the moles by the smallest amount of mole to get the empirical formula
C: 0.6mol C/0.6mol C = 1
H: 1.2mol H/0.6mol C = 2
O: 0.6mol O/0.6 molC = 1
Therefore, the empirical formula for this compound would be: CH2O
For further examples, here is a video to show you the some conversions between empirical and molecular formulas.
According to the periodic table, 1 mole of Sn = 118.7g 1.21ｘ10⁻⁸Sn moles ｘ 118.7g / 1mole = 143.627 g of Sn So, the mass of 7.32ｘ10¹⁵ Tin atoms is 144 grams.
Converting GRAMS into PARTICLES Ex. How many atoms of Hydrogen are present in 23.0g of NH₃?
Convert GRAMSinto MOLES
Since 1mole of NH₃= 17.0g, 23.0g of NH₃ ｘ 1mole / 17.0g
= 1.35 moles of NH₃
2. Then, convert MOLES into PARTICLES
1.35 moles of NH₃ｘ 6.022ｘ10²³ molecules/ 1mole = 8.13ｘ10²³ molecules of NH₃ 8.13ｘ10²³ molecules of NH₃ｘ 3H atoms / 1 molecule NH₃
= 2.44ｘ10²⁴H atoms
So, 2.44ｘ10²⁴atoms of Hydrogen are present in 23.0g of NH₃.
Here are some mole jokes that you might enjoy~!
Why are moles bad at counting?
Because they only know one number. What's the mole's favorite brand of soda? Coca-Mola.
Why are moles always on the phone? Because they love moleble devices. Why do moles love Tyra Banks? Because she's on America's Next Top MoledelWho is the the mole's favorite actor? Mole Gibson
What does Avogadro put in his hot chocolate? Marsh-mole-ows!
The above diagrams shows The conversion steps between Grams, Moles and Particles.
Now we are going to try to apply these:
(Remember sig figs!!)
ex/ How many moles of CCl4 are present in 8.36 x 10^35 molecules of CCl4?
Particle/Atom/Formula Unit ---> Moles
8.36 x 10^35 x 1/6.022 x 10^23 = 1.39 x 10^12 moles
ex/ How many copper atoms are present in 3.1 x 10^-3 mole of Cu2SO4
Moles ---> Particle/Atom/Formula Unit
Here it gets a little tricky, you are asked to find the number of copper atom, but it is Cu2, therefore you're gonna have to times 2 to get your final result.
3.1 x 10^-3 x 6.022 x 10^23 x 2 = 3.7 x 10^21 copper atoms
ex/ What is the mass in grams of 6.663 moles of neon gas?
Moles ---> Grams
Molecular mass of Ne2 = 2 Ne x 20.2 = 40.4 u
Molar mass of Ne2 = 40.4 g/mole
6.663 x 40.4 g/mole = 269 g Ne2
ex/ How many moles are there in 92.0 grams of Lead?
Mole is a unit like milometers, seconds, or litres.
A long time ago, scientists studied different gases and determined their properties.
They found the masses of hydrogen, oxygen, and carbon dioxide.
By finding the mass, they also discovered that if the gases had the same volume, they have the same ratios
Oxygen : Hydrogen = 16:1
Carbon Dioxide : Hydrogen = 22:1
Carbon Dioxide : Oxygen = 11:8
Different Types of Masses:
Relative Mass: a comparison (mathematically) between the mass of two objects
Hydrogen and Oxygen were all previously used as a standard for comparison but now, we use CARBON (mass= 12amu)as the standard for comparison.
Avogadro is a scientist that recognized the pattern and created:
AVOGADRO'S HYPOTHESIS
Equal volumes of different gases at the same temperature and pressure have he same number of particles
If the number of particles are the same, that means the mass ratio is due to the mass of the particles.
This principle was then used to create the periodic table.
Atomic Mass: mass of 1 atom of an element in atomic mass units
Example: Sodium has an atomic mass of 23.0 amu
Formula Mass: Mass of all the atoms in an ionic compound
Example: NaCl (salt)
Na = 23.0 amu
Cl = 35.5 amu
23.0 + 35.5 = 58.5 amu
Molecular Mass: Mass of all the atoms in a covalent compound
Example: CO
Carbon = 12.0 amu
Oxygen = 16.0 amu
12.0 + 16.0 = 28.0 amu
Molar Mass: Mass of one mole of each element
Example: 1 mole of oxygen = 16.0g/mol
1 mole of fluorine = 19.0 g/mol
Avogadro's number : 6.022x10^23
This is the number of particles in 1 mole of a substance.
The mole is just a unit like kilometers and meters and when it is used to express the molar mass of an atom, the unit for mole is grams/mol.
The mole is important because it allows chemists to use their mass to calculate how much atoms and molecules there are instead of actually counting them.
To give you an idea of how big one mole is, here is a video.
Objectives: To calculate the thickness of a sheet of aluminum foil and express the answer of proper scientific notation and significant figures.
Equipment
3 rectangular pieces of aluminum foil ( minimum 15 cm x 15 cm )
metric ruler
centigram balance
Procedure
Put the rectangular pieces of aluminum foil on each side and write number 1-3 on each respectively. Measure the length and width of each piece with a metric ruler. Record measurements ( using sig figs ) on table 1. Find the mass of each piece with a centigram balance and record it. Finally, compare and discuss results with groupmates.
Sheet Length ( cm ) Width ( cm ) Mass ( g )
1 15.5 14.6 0.88
2 16.7 15.1 1.05
3 14.5 14.9 1.02
The thickness of each foil 2.7 = density ( D = M/V )
1 ) 2.7 = 0.88 / V 2.38 = 15.5 x 14.6 x h
V = 2.38 cm cubed h = 1.05 x 10^-2
2 ) 2.7 = 1.05 / V 0.39 = 16.7 x 15.1 x h
V = 0.39 cm cubed h = 1.55 x 10^-3
3 ) 2.7 = 1.02 / V 0.38 = 14.5 x 14.9 x h
V = 0.38 h = 1.8 x 10^-3
Average thickness : ( 1.05 x 10^-2 + 1.5 x 10^-3 + 1.8 x 10^-3) / 3
= 4.6 x 10^-3
Accepted value = 1.55 x 10^-3
* sheet 2 has the most accurate result.
Here is a QUIZ to help you further understand mass volume and density and how to apply them to questions.
