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Chapter 10

Chapter 10. Chemical Quantities. Before We Begin…. I can write numbers in scientific notation. I can write numbers in standard notation. I can multiply numbers written in scientific notation. I can divide numbers written in scientific notation. Before We Begin….

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Chapter 10

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  1. Chapter 10 Chemical Quantities

  2. Before We Begin… • I can write numbers in scientific notation. • I can write numbers in standard notation. • I can multiply numbers written in scientific notation. • I can divide numbers written in scientific notation.

  3. Before We Begin… • We need to review some scientific notation. • Scientific notation is a way of writing very large and very small numbers.

  4. How to Write Numbers in Scientific Notation • Always written as a coefficient multiplied by 10 raised to a power. 3.5 x 1034 coefficient power

  5. Examples: • Write the following in scientific notation: • 234560000 • 0.00056974 • 8524000000 • 0.000000044258

  6. How to Multiply in Scientific Notation • To multiply numbers written in scientific notation you multiply the coefficients and add the powers. (2.35x1014) x (3.25x10-23) Multiply Add (2.35x3.25) x 1014+-23

  7. How to Multiply in Scientific Notation • To multiply numbers written in scientific notation you multiply the coefficients and add the powers. (2.35x1014) x (3.25x10-23) Multiply Add Answer = 7.64x10-9

  8. Examples: • Multiply the following numbers: • (1.23x104) x (4.56x107) • (7.89x10-1) x (1.23x1010) • (4.56x107) x (7.89x10-10) • (1.23x10-11) x (4.56x10-23)

  9. How to Divide in Scientific Notation • To divide numbers written in scientific notation you divide the coefficients and subtract the powers. (2.35x1014) ÷ (3.25x10-23) Divide Subtract

  10. How to Divide in Scientific Notation • To divide numbers written in scientific notation you divide the coefficients and subtract the powers. (2.35x1014) ÷ (3.25x10-23) Divide Subtract Answer =0.72x1037

  11. Examples: • Divide the following numbers: • (1.23x104) ÷ (4.56x107) • (7.89x10-1) ÷ (1.23x1010) • (4.56x107) ÷ (7.89x10-10) • (1.23x10-11) ÷ (4.56x10-23)

  12. Section 1 The Mole: A Measurement of Matter

  13. Section 1 Learning Targets 10.1.1 – I can describe methods of measuring the amount of something. 10.1.2 – I can define Avogadro’s number as it relates to a mole of a substance. 10.1.3 – I can distinguish between the atomic mass of an element and its molar mass. 10.1.4 – I can describe how the mass of a mole of a compound is calculated.

  14. Measuring Matter • You often measure the amount of something by one of three different methods – by count, by mass, and by volume.

  15. Example: • If 0.20 bushel is 1 dozen apples and a dozen apples has a mass of 2.0kg, what is the mass of 0.50 bushel of apples?

  16. What Is a Mole? • Mole (mol) – 6.02x1023 representative particles of that substance (SI unit for measuring the amount of something). • Avogadro’s number - 6.02x1023 named after Amadeo Avogadro diQuarenga (1776-1856)

  17. A mole of any substance contains Avogadro’s number of representative particles, or 6.02x1023 representative particles.

  18. Converting Number of Particles to Moles • You can use Avogadro’s number as a conversion factor.

  19. Example: • How many moles is 2.80x1024 atoms of silicon?

  20. Converting Moles to Number of Particles • The reverse also works.

  21. Example: • How many molecules are in 5.6 moles of NO2?

  22. The Mass of a Mole of an Element • The atomic mass of an element expressed in grams is the mass of a mole of the element. • Molar mass – the mass of a mole of an element. • Find the element on the periodic table and the mass that’s listed is the mass of one mole.

  23. The Mass of a Mole of a Compound • To calculate the molar mass of a 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.

  24. Example: • What is the mass of 1.00 mol of sodium hydrogen carbonate?

  25. Section 2 Mole-Mass and Mole-Volume Relationships

  26. Section 2 – Learning Targets 10.2.1 – I can describe how to convert the mass of a substance to the number of moles of a substance, and moles to mass. 10.2.2 – I can identify the volume of a quantity of gas at STP.

  27. The Mole-Mass Relationship • Use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance.

  28. Example: • Find the mass, in grams, of 4.52x10-3mol of C20H42.

  29. The reverse is also true.

  30. Example: • Calculate the number of moles in 75.0g of dinitrogen trioxide.

  31. The Mole-Volume Relationship • Avogadro’s hypothesis – states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles.

  32. Standard temperature and pressure (STP) – means a temperature of 0°C and a pressure of 101.3kPa or 1 atmosphere (atm).

  33. At STP, 1 mole or 6.02x1023 representative particles, of any gas occupies a volume of 22.4L • Molar volume – the 22.4L of a gas.

  34. Calculating Volume at STP • 22.4L = 1 mol at STP provides a nice conversion factor.

  35. Example: • What is the volume of 3.70 mole N2 at STP?

  36. Example • How many moles are in 102 L of carbon dioxide, CO2?

  37. Calculating Molar Mass from Density • Different gases have different densities and is usually measured in g/L so we can calculate different things using density as a conversion factor.

  38. Example: • A gaseous compound composed of sulfur and oxygen, which is linked to the formation of acid rain, has a density of 3.58 g/L at STP. What is the molar mass of this gas?

  39. The Mole Road Map • A helpful tool to figure out easily which conversion factor to use. This can also be found on page 303 in your Chemistry book

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