1 / 36

The Mole

The Mole. Chemical formulas are not simply abbreviations for words. They represent precise quantities. What 2 “bits” of information do formulas give? Water for example is H 2 O. It indicates that 1 molecule of water consists of 2 hydrogen atoms and 1 oxygen atom.

ipo
Download Presentation

The Mole

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. The Mole Chemical formulas are not simply abbreviations for words. They represent precise quantities. What 2 “bits” of information do formulas give? Water for example is H2O. It indicates that 1 molecule of water consists of 2 hydrogen atoms and 1 oxygen atom. How could you represent 2 moleculesof water? Use a coefficient This formula shows 4 hydrogen atoms and 2 oxygen atoms. • Elements • Number of atoms of each element 2 H2O

  2. In daily life, pieces of matter are often measured either by counting them or by massing them; the choice is determined by convenience. If you buy eggs, you buy them by the dozen—that is, by number. Eggs are easy to count out. So are oranges and lemons. Other items, though countable, are more conveniently sold by mass. A dozen peanuts is too small a number to buy, and several hundred are too difficult to count. You buy peanuts by the pound or kilogram—that is, by mass. Chemists, too, are interested in quantity—the quantity of an element or compound, which, like grocery items, can be measured by number or by mass. Although you can easily mass a sample of a substance, yet the number of atoms or molecules in it is much too large to count. Nevertheless, chemists are interested in knowing such numbers.

  3. In the laboratory, when we “run” a reaction we want to know the number of atoms, molecules, or formula units in a substance because these are the entities that react with each other. However, these entities are much too small to count individually, so chemists use a unit called the mole to count by massing them. • The mole was derived from the Latin word moles, meaning “heap” or “pile.” The mole, whose symbol, mol, is the SI base unit for measuring the amount of substance. • A mole is the amount of substance that is equal to the number of carbon atoms in exactly 12 g of the carbon-12 isotope. • One mole always contains the same number of particles, no matter what the substance. 602,200,000,000,000,000,000,000 • 1 mole = 6.022 x 1023 particles This value is referred to as Avogadro’s number (Amadeo Avogadro) in honor of the man who conceived the basic idea.

  4. One mole each of various substances - Clockwise from top left: 1-octanol (C8H17OH); mercury(II) iodide (HgI2); methanol (CH3OH); sulfur (S8).

  5. The central relationship between the mass of one atom and the mass of one mole of atoms is that the mass of an element is expressed in amu. The mass of one mole of atoms of an element is the same numeric value but measured in grams. 1 atom of sulfur has a mass of 32.07amu. 1 mole of sulfur atoms has a mass of 32.07 g.

  6. So, • 1 mole S = 6.022 x 1023 atoms S = 32.07 g S • 1 mole of Fe2+ = 6.022 x 1023 ions = 55.85 g Fe2+ • 1 mole H2O = 6.022 x 1023 molecules H2O = 18.02 g H2O • 1 mole NaCl = 6.022 x 1023 formula units NaCl = 58.44 g NaCl • How can I get the mass of one unit of a compound? • Add up the individual masses of theelements. • If we are talking about themass of an atomwe call it the atomic mass. • If a compound isa molecule, we call this themolecular mass. • If a compound isa formula unit, we call it theformula mass.

  7. g-formula mass = 1 MOLE = 6.022 x 1023 “particles” • 1 mole • of asubstance = Avogadro’s constant = (6.022 x 1023) • mass of substance • in grams elementschargedcovalentionic particlescompoundscompounds Atoms Ions Molecules Formula Units

  8. The Mole Match the representative particle with the appropriate substance. CCl4 lead (IV) bromide Rb carbon monoxide Sr2+ Ba(NO3)2 H2 CO32- tungsten Al(CN)3 How many potassium atoms are contained in one mole of the element? What is the mass of one mole of potassium? Atom Molecule Ion Formula unit 39.10 g 6.022 x 1023

  9. A 1-carat diamond has a mass of 0.200 g. How manymolesof carbon? How manyatomsof carbon? HINT: If moles of your substance does not appear in your given or in the desired units, then you must do at least2 conversion factors. • USE FLM!!!!!!! • Start with your given, identify units ofanswer, and write conversion factors.

  10. A ring is constructed out of 3.06 x 1022 atoms Au. How many grams? How many moles?

  11. Counting Atoms and Molecules • A raindrop contains • about 0.050 g of water. • How many molecules of water? • How many moles of O? • How many atoms of H? 1 molecule of H2O 1 mole of H2O 2 atoms of H 2 moles of H 1 atom of O 1 mole of O

  12. How many molecules of water? = 1.7 x 1021 molecules H2O • How many moles of O? = 0.0028 moles O

  13. YOU TRY THIS ONE How many atoms of H? or = 3.3 x 1021 atoms H

  14. Moles of What?!?! One teaspoon of table sugar (sucrose) is approximately 15 g of C12H22O11. How many moles of sugar is this? 0.044 mole C12H22O11 How many atoms of carbon are in this 15 g sample? = 3.2 x 1023 atoms C

  15. Moles of What?!?! What mass of CaCl2 would be needed to furnish 1.00 mole of Cl1- ions? What are we looking for as an answer?!?! = 55.5 g CaCl2

  16. Percent Composition Compare the mass of each element present in one mole of a compound to the total massof one mole of a compound. EXAMPLE:Find the % composition of Al2(SO4)3 • = 342.17 g Al2(SO4)3

  17. Percent Composition EXAMPLE:Find the % composition of Al2(SO4)3 Check the sum of the %’s. They should sum to very near

  18. EXAMPLE:Find the mass percent (% composition) of Ferrous nitrate Fe(NO3)2 What is the formula for ferrous nitrate? = 179.87 g Fe(NO3)2

  19. If you had a 100.0 g sample of ferrous nitrate, how many grams of iron are there? 31.05 g Fe 15.58 g N Nitrogen? If you had a 540.0 g sample of ferrous nitrate, how many grams of iron are there? 540.0 x (.3105 % Fe) = 167.7 g Fe Nitrogen? 540.0 x (.1558 % N) = 84.13 g N

  20. Which toothpaste component has a higher % by mass of fluoride? CaF2 NaF CaF2 Calcium fluoride 48.67% F, while sodium fluoride 45.25% F

  21. Empirical Formulas We can also do the reverse and use percent composition data to determine empirical and molecular formulae. Empirical formulas simplest whole number ratio of atoms in a compound. Molecular formulas show the actual numbers of atoms of each element in a compound.

  22. Remember the tune: Percent to mass Mass to mole Divide by the smallest Multiply ‘til whole

  23. There are two ways to determine empirical formulas. Experimental Data% composition data Need to know mass of Assume the amount • each element in the to be 100 g laboratory sample

  24. EXAMPLE: 40.0% C 6.71% H 53.3% O 1. Assume 100 g sample 40.0 g C 6.71 g H 53.3 g O 2. Convert to moles

  25. Find the mole ratio • (divide by the smallest number of moles.) If the mole ratio ends in .20 (x5) .25 (x4) .33 or .66 (x3) .50 (2) Otherwise round to one of the above decimals or to the nearest whole number. 4. Empirical Formula CH2O or CHOH

  26. EXAMPLE: 37.70% Na, 22.95% Si, 39.34% O Na2SiO3

  27. EXAMPLE: A binary compound consisting only of nitrogen and oxygen with a mass percent of 37.00% N

  28. multiply by 2 to eliminate and make whole number ratios. N2O3

  29. Molecular Formulas Molecular formulas show the actual number of atoms of each element in a molecule, as well as, the ratio of atoms. In order to solve these types of problems, 3 pieces of information are needed. • molecular mass of the compound • empirical mass of the compound • number of empirical units present in the molecular formula

  30. EXAMPLE: What is the formula of an unknown substance with a percent composition of 5.90% H and 94.10% O. It also has a molecular mass of 34.00 g/mol. 5.90% H and 94.10% O means 5.90 g H and 94.10 g O

  31. Empirical Formula is:HO The empirical mass is 17.01 g (1.01 + 16.00) We need to determine how much bigger the molecular formula is compared to the empirical. Molecular mass is given in the problem Molecular formula is 2 times the empirical • Empirical Formula is: HO • Molecular Formula is: H2O2

  32. Hydrates Hydrates are crystals that contain water molecules in their crystal structure. The crystal has crystallized from a water solution with molecules of water adhering to the particles of the crystal. EXAMPLE: NiSO3. 6 H2O 1 crystal of NiSO3 with 6 molecules of H2O

  33. To calculate the formula of a hydrate, we need: • mass of the hydrated sample • mass of dry sample (anhydrous, w/o water) after the H2O has been driven off. • Mole ratio of the anhydrous compound to the H2O driven off. • Finding the formula of a hydrate is very similar to finding empirical formulas.

  34. EXAMPLE: We have a 10.407 g sample of hydrated barium iodide. The sample is heated to drive off the water. The dry sample has a mass of 9.520 g. What is the formula of the hydrate? The difference between the initial mass and that of the dry sample is the mass of water that was driven off.

  35. Formula: BaI2. 2 H2O

More Related