1 / 26

Mastering Moles and Molar Mass: A Comprehensive Guide

Dive into the world of moles, empirical and molecular formulas, and hydrates in this detailed guide. Discover the concept of a mole, Avogadro’s Number, molar mass calculations, and mole conversions. Learn how to convert between grams and moles, particles and moles, and particles and grams using Avogadro’s Number and molar mass. Practice problems and examples included for easy understanding and application. Explore empirical formulas, percent composition, and how to determine the empirical formula of a compound. Enhance your chemistry knowledge today!

guillermoh
Download Presentation

Mastering Moles and Molar Mass: A Comprehensive Guide

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 10: Moles, Empirical & Molecular Formulas, Hydrates

  2. What is a Mole? SI unit for amount of substance is called mole A mole measures the number of particles/atoms/molecules within a substance A mole refers to a specific number of particles (it is a counting unit, similar to a dozen)

  3. 1 mole = 6.02 x 1023 particles/atoms/molecules 6.02 x 1023 is also known as Avogadro’s Number 1 mole magnesium = 6.02 x 1023Mg atoms 1 mole zinc = 6.02 x 1023Zn atoms 1 mole lead = 6.02 x 1023Pb atoms

  4. Molar Mass Although 1 mole always contains the same number of particles(6.02 x 1023), the mass of one mole varies depending on the substance Molar Mass: The mass of onemole of a substance. Mass of one mole of an element is EQUAL to its ATOMIC MASS,which isexpressed in grams. 1 mole of magnesium = 24.31 grams 1 mole of zinc = 65.39 grams 1 mole of lead = 207.2 grams

  5. Molar Mass of a Compound • To calculate the molar mass a compound, first determine the molar mass of each element in the compound • Then, add together each element’s molar mass • The sum is the molar mass of the compound • Example #1: Determine the molar mass of NaOH. • Example #2:Determine the molar mass of HCN.

  6. Mole Conversions • Many times in chemistry, you will need to express a measurement in a unit different from the one given or measured initially • Conversion Factor:Ratio of equivalent measurements used to convert from one unit to another

  7. Mole Grams Molar Mass Mole ⬄ Grams Conversions • To convert between grams and moles, use molar mass as the conversion factor Example #1:mol⬄ grams — How many grams are in 3.01 mol of iron (Fe)? Example #2:grams ⬄ mol — How many moles are in 75 g of Fe? Example #3: mol⬄ grams — What is the mass of 2.55 mol of the compound KMnO4? Example #4:grams ⬄ mol – How many moles are in 35.0 g of HCl?

  8. Mole ⬄ Grams Conversions: Example Problem Solutions Example #1: mol ⬄ grams — How many grams are in 3.01 mol of iron (Fe)? Example #2:grams ⬄ mol — How many moles are in 75 g of Fe? 75 g Fe x 1 mol Fe = 1.3 mol Fe 55.847 g Fe

  9. Mole ⬄ Grams Conversions: Example Problem Solutions Example #3: mol⬄ grams — What is the mass of 2.55 mol of the compound KMnO4? 2.55 mol x 158 g = 402.9 g x 1 mol Example #4:grams ⬄ mol – How many moles are in 35.0 g of HCl? 35.0 g x 1 mol = 0.959 mol x 36.5 g

  10. Mole ⬄ Particles/Atoms Conversions • To convert between particles/atoms and moles, use Avogadro’s Number (6.02 × 1023) Example #1:particles ⬄ mol — How many moles equal 2.41 × 1024 atoms of aluminum (Al)? Example #2:mol ⬄ particles — How many atoms are in 3.45 mol of Al? Particles Mole Avogadro’s #

  11. Mole ⬄ Particles/Atoms Conversions: Example Problem Solutions Example #1:particles ⬄ mol — How many moles equal 2.41 × 1024 atoms of aluminum (Al)? Example #2:mol ⬄ particles — How many atoms are in 3.45 mol of Al?

  12. Mole Conversion Practice, (cont.) 2.) A chemist produced 11.25 g of magnesium (Mg). How many moles of Mg were produced? 3.) Determine the number of silver (Ag), atoms that are contained in 0.650 mol of Ag. 11.25 g x 1 mol = 0.463 mol x 24.3 g 0.65 mol x 6.02 x 1023 atoms = 3.91 x 1023 atoms x 1 mol

  13. Particle Mole Avogadro’s # Molar Mass Gram Particles ⬄ Grams Conversions • To convert between particles and grams, use BOTHmolar mass ANDAvogadro’s Number (cannot go directly from particles to grams and vice versa; must first find MOLES) Example #1:particles ⬄ grams — How many grams are in 7.00 × 1033 molecules of H2O? Example #2:grams ⬄ particles — How many molecules are in 8.25 g of H2?

  14. Particles ⬄ Grams Conversions: Example Problem Solutions Example #1:particles ⬄ grams — How many grams are in 7.00 × 1033 molecules of H2O? Example #2:grams ⬄ particles — How many molecules are in 8.25 g of H2?

  15. Empirical Formulas

  16. Percent Composition • The percent by mass of an element in a compound • Percent by mass (element) = mass of element x 100 • mass of compound • For example, in H2O, you have 2 g of hydrogen and 16 g of oxygen. The total mass of the compound is 18 g. • Hydrogen: 2 g x 100 = 11% hydrogen • 18 g • Oxygen:16 g x 100 = 89% oxygen • 18 g • The percentages should always equal 100%

  17. Percent Composition • Practice Problem • Sodium bicarbonate (NaHCO3), commonly known as baking soda, is an active ingredient in some antacids used for the relief of indigestion. Determine the percent composition of each element in NaHCO3. • NaHCO3 formula mass = 84 g • Elements = Na, H, C, O

  18. Determining the Empirical Formula of a Compound If we know how much of each element is present in a compound, we can determine the formula of the compound Amounts may be expressed in: percentages grams moles If composition is given in terms of percentages, assume you have a 100 g sample of the compound

  19. The empirical formula, or simple formula, is a formula in which the elements are in their lowest whole number ratio. The empirical formula may or may not be the actual formula. Example: The formula for glucose is C6H12O6 . Its empirical formula is CH2O. The subscripts in a chemical formula represent a MOLE ratio of elements in the compound. If given mass (grams), must convert to moles

  20. DETERMINING THE EMIRICAL FORMULA OF A COMPOUND Example #1: Analysis shows a compound consists of 32.4% Na, 22.6% S, and 45.0% O. Determine the empirical formula for this compound. Example #2: A compound is found to contain 4.43 g phosphorus and 5.72 g oxygen. Determine the empirical formula for this compound.

  21. Molecular Formulas • Molecular Formula: Actual formula of a compound; may or may not be the same as the empirical formula • Actual formula is a multiple of the empirical formula. • C6H12O6 = 6 (CH2O) • In order to find the molecular formula, you must be given the formula mass of the compound. • Divide the formula mass by the empirical mass. This will give you the multiple. • CH2O = 30 g • C6H12O6 = 180 g In this case, the multiple is 6.

  22. DETERMINING THE MOLECULAR FORMULA OF A COMPOUND • Example #1: Determine the molecular formula of a compound with an empirical formula of CH and a formula mass of 78 amu. • Example #2: A compound with a formula mass of 60.0 amu is found to be 39.9% carbon, 6.7% hydrogen, and 53.4% oxygen by mass. Find its molecular formula.

  23. Hydrates

  24. Hydrates • Hydrate: A compound that has a specific number of water molecules bound to its atoms • Formula for a hydrate is written as follows: • Formula for ionic compound x H2O • x = number of water molecules

  25. Molar Mass of Hydrates • When given the formula of a hydrate, the water molecules are included in the molar mass along with the ionic compound • Example #1: CuSO4 5 H2O • CuSO4 = 159.6 g • H2O = (5) 18.0 = 90 g • 159.6 g + 90 g = 249.6 g • Example #2:BaCl2 2 H2O • BaCl2 = 208.3 g • H2O = 2 (18.0) = 36 g • 208. 3 g + 36 g = 244.3 g

  26. Hydrates

More Related