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Using Chemical Formulas

Using Chemical Formulas. Bring your calculators to class. Remember the mole? (not just a furry animal that digs holes in the yard.) unit used by chemist to measure things.

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Using Chemical Formulas

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  1. Using Chemical Formulas Bring your calculators to class

  2. Remember the mole?(not just a furry animal that digs holes in the yard.) • unit used by chemist to measure things. • 1 mole is 6.02 x 1023 particles (like a very large dozen used for very small things) • Defined as the number of carbon atoms in exactly 12 grams of carbon-12. • Makes life very convenient

  3. What sort of things are measured in moles? Representative particles(The smallest pieces of a substance.) For a molecular compound it is a molecule. H2O For an ionic compound it is a formula unit. NaCl For an element it is an atom. Au - - + + - - + +

  4. Molar Mass generic term for the mass of one mole. Same as: Don’t Write 1.gram atomic mass =The mass of 1 mole element ingrams 2.gram formula mass =the mass of 1 mole of one formula unit is the sum of the atomic masses of all atoms of element 3.gram molecular mass =the mass of 1 mole of one molecule is the sum of all the atomic masses of all the atoms of elements in the molecule

  5. 1. Elements The mass of 1 mole element in grams 1 mole of any element = atomic mass on periodic table example: 1 mole of Na=23 g (22.989768 rounded to whole #) 1 mole of Cl = 35g (35.4527 rounded to 35)

  6. 2. Ionic compounds the mass of 1 mole of one formula unit is the sum of the atomic masses of all atoms of element example: 1 mole of NaCl= 23 mass of Na + 35 mass of Cl 58 g

  7. 3. Molecular compounds the mass of 1 mole of molecules is the sum of all the atomic masses of all the atoms of elements in a molecule example: 1 mole of H2O2= 2 x 1g = 2g mass of H 2x16 + 32g mass of O 34g

  8. Examples Calculate the molar mass of the following and tell what type it is. Na2S N2O4 C Ca(NO3)2 C6H12O6

  9. Formula Mass = Molar Mass Sum of the atomic masses of all the atoms in a chemical formula H2O 2 H atoms each 1g= 2x1= 2 1 O atom 16 g +16 18 g 1mole of H2O = 18 g 1 mole = 6.022 x 1023 particles

  10. Using Molar Mass Molar Mass =The number of grams of 1 mole of atoms, formula units, or molecules. We can make conversion factors from these. To change grams of a compound to moles or moles to grams of a compound use 1 mole= molar mass to make factor

  11. For example: mass to moles How many moles is 5.69 g of NaOH? need to change grams to moles Molar mass for NaOH 1mole Na = 23g +1 mole O = 16g +1 mole H = 1g 1 mole NaOH = 40g 5.69g NaOH 1mole NaOH 40 g NaOH = 0.14 mole

  12. Another Type: moles to mass Molar mass of H2O 2 x 1g = 2 g of H +16g of O 18g 2 moles H2O 18 g H2O 1 mole H2O =36 g H2O How many grams are 2 moles H2O? Need to change from moles to grams 1 mole of H2O= 18g

  13. Other Types of questions How many molecules of CO2 are the in 4.56 moles of CO2 ? 1 mole = 6.02 x 1023 particles 4.5 moles 6.02 x 1023 molecules = 1 mole 2.79 x 1024 molecules

  14. 7.78 x 1024formula units 1 mole 6.02 x 1023 formula units 1.29 x 10 or 12.9 moles How many moles is 7.78 x 1024 formula units of MgCl2?

  15. Examples How much would 2.34 moles of carbon weigh? How many moles of magnesium in 24.31 g of Mg? How many atoms of lithium in 1.00 g of Li? How much would 3.45 x 1022 atoms of U weigh?

  16. Hydrates- when some salts crystallize from a water solution and they bind water molecules in their crystal structure Cu SO4•5H2O there are 5 water molecules for every copper(II) sulfate formula unit copper(II) sulfate pentahydrate

  17. Heating the crystal in a crucible drives off the water then the salt is called anhydrous CuSO4•5H2OCuSO4 + 5H2O

  18. CuCl2•2H2O CuCl2 + 2H2O

  19. % Composition Like all percents part x 100 % = % whole 1. Find the mass of each component 2. divide by the total mass x 100 % Example: Calculate the percent composition of Na in NaCl? mass Na x 100 % = 23 g x 100 % =39% mass NaCl 23 + 35

  20. Empirical Formulas -simplest whole number ratio of atoms % to mass mass to mole divide by small multiply ‘til whole

  21. Example What is the empirical formula of the compound that contains 53.73 % Fe and 46.27 % S? 1. % to mass assume 100 g so % = mass 53.73 g Fe & 46.27 g S

  22. 2. Mass to mole 1 mole Fe = 55.85 g 53.73 g Fe 1mole = .96 mole Fe 55.85 g Fe 1 mole S = 32.06 g 46.2 g S 1 mole S = 1.44 mole S 32.06 g S

  23. 3. Divide by small Fe .96/.96 = 1 S 1.44/.96 = 1.5 4. Multiply ‘til whole Fe 1 x 2 = 2 S 1.5 x 2 = 3 Write empirical formula Fe2S3

  24. Molecular formula = How the molecule actually exists If the molar mass of the compound is known, then the molecular formula can be determined from the empirical formula. 1. Add the masses of all the atoms of each elements in the empiricalformula. 2.Divide the molar mass by the mass determined in step 1. 3.Multiply each subscript in the empirical formula by the number calculated in step 2.

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