1. Measuring Matter
2. The mole: 6.02 x 10²³ [ Practice] [Reference]
3. The gram formula mass
4. The molar mass of a substance [Practice ]
5. Mole-mass conversions [Practice ]
6. The volume of a mole of gas [Practice ]
7. Gas density and gram molecular mass [Practice ]
8. Converting between units with moles [Practice ]
9. Calculating % composition [Practice ]
10. Calculating empirical formulae [Practice ]
11. Calculating molecular formulae [Practice ]
12. Scientific Notation [Practice 1 ] [Practice 2]

Chapter 6

6-1 Measuring Matter

- Counting something is one way of measuring something.

- The SI unit for the amount of something is the mole.

6-2 The Mole [Practice ]

- A mole of something is Avogadro's number of that thing. Avogadro's number is 6.02 x 10^23. A mole of pennies is 6.02 x 10^23 pennies.

6-3 The Gram Formula Mass

- Gram atomic mass: same as atomic mass, but using grams as units instead of AMU's. Carbon's atomic mass is 12.0 AMU's. Carbon's gram atomic mass is 12.0 g.

- The gram atomic mass of an element contains a mole of that element's atoms.

- Gram molecular mass: add AMU's of all atoms in one molecule, substitute "grams" for "AMU's."

- Gram formula mass: add AMU's of all atoms in one formula unit, substitute "grams" for "AMU's."

6-4 The Molar Mass of a Substance [Practice]

- The molar mass of something is the mass of a mole of it. It is identical to the element's gram formula mass.

6-5 Mole-Mass Conversions [ Practice]

- Mole-Mass conversions can be made using the true relationships method described in section 3-4.

6-6 The Volume of a Mole of Gas [Practice]

- Standard Temperature and Pressure (STP): 273 K, 1 atm.

- 1 mol of any gas at STP occupies 22.4 liters.

6-7 Gas Density and the Gram Molecular Mass [Practice]

- True relationships (section 3-4) can be used to find gram molecular mass given density.

6-8 Converting Between Units and Moles [Practice]

- Now that we have learned so many different types of units, converting between them can get complicated. The mole makes conversions much simpler if you know how to use it. Chempire's Online Conversion Map can help out greatly in this area.

6-9 Calculating Percent Composition [Practice]

- Percent composition describes what percentage, by mass, an element takes up in a compound.

- % mass = (grams of element)/(grams of compound) x 100%.

6-10 Calculating Empirical Formulas [Practice]

- A compound's empirical formula is the lowest whole-number ratio of elements in that compound.

- Example:

 - We have a compound that is 25.9% nitrogen and 74.1% oxygen. What is the compound's empirical formula.

 - In 100g of the compound, there is 25.9g N and 74.1g O:

 25.9 g N  x    1 mol N  =  1.85 mol N
                     14.0 g N

 74.1 g O  x  1 mol O   =  4.63 mol O
                   16.0 g O

 - Now we know the ratio of the elements: N(1.85)O(4.63). But we know that you can't have partial atoms, so we divide both numbers in the ratio by the  lowest one.

 - Now we have a better ratio NO(2.5). But still we have a partial atom in there. So we multiply both numbers in the ratio by 2 and we have an empirical  formula: N2O5.

6-11 Calculating Molecular Formulas [Practice]

- A molecular formula is either the same as the empirical one, or a multiple of it.

- Given the gram formula mass of a molecular compound and its empirical formula, you can find its molecular formula by dividing the gram formula mass by the empirical formula mass and using this answer to multiply all of the subscripts by.

6-12 Scientific Notation [Practice 1] [Practice 2]

- Always in form: #.### . . . x 10^x.

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