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Chemical Quantities

Chapter 10 explores the concept of the mole, a fundamental unit in chemistry that quantifies the number of particles in a substance. It defines a mole as 6.02 x 10²³ representative particles and explains its significance in chemical calculations. The chapter covers the relationships between moles, mass, and volume, including how to calculate molar mass and volume at standard temperature and pressure (STP). It also addresses percent composition, empirical formulas, and molecular formulas, providing a comprehensive understanding of chemical quantities in a quantitative world.

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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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