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Chemical Quantities

Chemical Quantities. Chemistry Chapter 10. The Mole: A Measure of Matter. We live in a quantitative world! How much or how many? We often measure the amount of something by one of three different methods – by count, by mass, and volume .

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Chemical Quantities

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  1. Chemical Quantities Chemistry Chapter 10

  2. The Mole: A Measure of Matter • We live in a quantitative world! • How much or how many? • We often measure the amount of something by one of three different methods – by count, by mass, and volume. • Must be able to convert from mass to volume to count, if given conversion factors.

  3. What is a mole? • A mole is a unit (used by chemist) to specify the number of particles. A mole (mol) of a substance is 6.02 x 1023 representative particles of that substance; an SI unit. • Representative particles refers to the species present in a substance: • Atoms – representative particle for most elements • Molecules – molecular compounds & diatomic molecules • Formula units – ionic compounds

  4. Avogadro’s Number • A mole of any substance contains Avogadro’s number (6.02 x 1023) of representative particles. • The mole as a conversion factor. • 1 mol = 6.02 x 1023 representative particles • Thus you can convert from # of particles to moles and from moles to # of particles. 1 mole_______ 6.02 x 1023 rep particles

  5. The Mass of a Mole • The atomic mass of an element expressed in grams is the mass of a mole of the element. • The mass of a mole of an element is itsmolar mass (g/mol). • How would you find the mass of a mole of an compound? • Find the number of grams of each element in one mole of the compound, then add the masses of the elements in the compound;formula mass (g/mol).

  6. The Mole – Mass Relationship • Use the molar (formula) mass of an element or compound to convert between the mass of a substance and the moles of a substance.

  7. The Mole – Volume Relationship • In 1811 Amedeo Avagadro proposed that equal volumes of gases (at the same temp and pressures) contain equal number of particles; known as Avogadro’s hypothesis. • Because of variations due to temp and pressure, the volume of a gas is usually measured at a standard temp and pressure (STP). STP means a temp of 0oC and a pressure of 101.3 kPa or atmosphere (atm). • At STP, 1 mole or 6.02 x 1023 rep particles, of any gas occupies a volume of 22.4 L (molar volume of a gas).

  8. The Mole – Volume Relationship • Calculating volume at STP – the molar volume is used to convert a known number of moles of gas to the volume of the gas at STP. Calculating Molar Mass from Density – the density of a gas is measured in grams per liter (g/L). The density of a gas at STP & the molar volume at STP (22.4 L/mol) can be used to calculate the molar mass of the gas.

  9. The Percent Composition of a Compound • Relative amounts of elements in a compound are expressed as the percent composition(or the percent by mass)of each element in the compound. • The percent by mass of an element in a compound is the number of grams of the element divided by the mass in grams of the compound, multiplied by 100%.

  10. Suppose you want to know how much carbon and hydrogen are contained in 82.0 g of propane ?

  11. In the previous problem we found that propane is 81.8% carbon and 18% hydrogen. That means that in a 100-g sample of propane, you would have 81.8 g of carbon and 18 g of hydrogen. You can use the ratio • To calculate the mass of carbon contained in 82.0 g of propane (C3H8). • Using the ratio 18 g H/100 g C3H8,you can calculate the mass of hydrogen The sum of the two masses equals 82 g, the sample size!!!

  12. Empirical Formulas • The empirical formula of a compound shows the smallest whole-number ratio of the atoms in the compound. • Is calculated using the percent compositions • Calculate % composition (will be given) • Change ratio of masses to moles, using molar mass conversion factor. • Reduce ratio to lowest whole number (these become the subscripts). • May or may not be the same as a molecular formula, which tells you the actual # of each kind of atom in a molecule.

  13. Molecular Formulas • A molecular formula of a compound is either the same as its experimentally determined empirical formula, or it is a simple whole-number multiple of its empirical formula. • Given the empirical formula you can determine the molecular formula; however you must have (be given) molar mass of the substance. • Calculate the EFM (empirical formula mass). Then divide the actual molar mass by the efm; multiply the subscripts in the empirical formula by this value to obtain the molecular formula.

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