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Chemical Quantities

Chemical Quantities. Standards. 3b. Students know the quantity of one mole is set by defining one mole of carbon-12 atoms to have a mass of exactly 12 grams. 3c. Students know one mole equals 6.02 x 10 23 particles (atoms or molecules).

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Chemical Quantities

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  1. Chemical Quantities

  2. Standards • 3b. Students know the quantity of one mole is set by defining one mole of carbon-12 atoms to have a mass of exactly 12 grams. • 3c. Students know one mole equals 6.02 x 1023 particles (atoms or molecules). • 3d. Students know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure.

  3. Objectives • Know the meaning of the mole • Know how to convert between moles, grams, and number or particles • Know how to calculate the molar mass of a compound • Know how to calculate percent composition • Know how to determine empirical, molecular and hydrate formulas

  4. Identity of some basic “particles” in chemistry • Element • Atom • Molecule…diatomic elements

  5. Compounds Covalent vs Ionic • Covalent Compounds • Are composed of molecules

  6. CompoundsCovalent vs Ionic • Ionic Compounds • Are composed of positive and negative ions; the representative particle is called a formula unit • Formula unit: the simplest ratio of ions represented in an ionic compound

  7. Mole • Defn: a counting reference. A quantity chosen by chemists to represent the amount of a substance. • It is the number of representative particles in exactly 12 grams of pure carbon-12 • The amount of a substance which contains “6.022 x 1023 particles” • Particles can be: Atoms, molecules, formula units, ions, etc. • 6.022 x 1023 is known as Avogadro’s Number • 1 mole = the atomic mass of that element in grams...from the PT

  8. Question • Where do you find the atomic mass for an element? • 12 grams of carbon contains the same # of atoms as 19 grams of fluorine or 78.96 g of Se because they each contain 1 mole of particles • 1 mole of He = ________g He • 58.69 g Ni = ________mole Ni

  9. Atomic Mass • Atomic Mass: the weighted average mass of the isotopes of an element…the value on the PT • What is the atomic mass of hydrogen? • 1.00794 g • What is the atomic mass of Oxygen? • 15.999 g

  10. Avogadro's Number • Avogadro's Number is 6.022x1023 which is the number of representative particles in a mole • 6.022x1023 “particles” = 1 mole

  11. Mole Conversions • Determine the “given” and the “get” • Mini road map • Set up equation with conversion factor(s) • Cancel the units and calculate. • Does your answer make sense? MOLES ↔ GRAMS MOLES ↔ ATOMS

  12. Example • 2.00 moles Boron = how many grams Boron?

  13. Example • How many moles of calcium are there in 23.20 grams of calcium?

  14. Example • How many atoms are there in 5.25 moles of nitrogen gas?

  15. Put the Moles in the Middle! • GRAMS ↔ MOLES ↔ ATOMS “MOLES IN THE MIDDLE,” you gotta go through the MOLE!

  16. Generic Mole Conversion Set-Up

  17. Example • How many atoms are there in 27.35 g of aluminum?

  18. Example • How many grams are there in 6.7 x 1024 molecules of HF?

  19. Molar Mass • Defn: the sum of the atomic masses of all the elements in a compound/element

  20. Calculating Molar Mass • Determine the atomic mass of each element in the compound • The atomic mass is found on the periodic table • Add all the atomic masses together, this is your molar mass • EX. What is the molar mass of water? H2O: 2 H = 2 ( 1.00794 g) 1 O = 15.9994 g/mol 2 ( 1.00794 g) + 15.9994 g= 18.0153g/mol

  21. Question • What is the molar mass of calcium sulfate?

  22. Percent Composition • The percent by mass of each element in a compound • To determine the percent composition of a compound: • Determine the molar mass of the compound • Determine the molar mass of the element • Determine the percent of each element in the compound, using the % composition formula • Check to make sure percentages add up to 100% • % El = (# atoms of El) (atomic mass El) x 100 molar mass of compound

  23. Question • What is the percent composition of H2O? • % El = (# atoms of El) (atomic mass El) x 100 molar mass of compound % H = (2) ( 1.00794 g) x 100 = 11.2% H 18.0153 g % O =? 100-11.2=88.8%

  24. Examples • Calculate the percent composition of the following compounds: C2H6 CO2 NaHSO4

  25. CO2

  26. NaHSO4

  27. Empirical Formulas • Defn: the formula with the smallest whole number mole ratio of elements in a compound • If given percent composition, then: • Assume a 100 gram sample • Find the number of moles of each element • Divide each mole number by the smallest mole number • Write the formula • If you come up with a decimal number, multiply the decimal number by a whole number to get a whole number • Then multiply all other ratios by that same whole number

  28. Example • Given: 94.11% O and 5.89% H Assume 100 g sample→

  29. Example • Given: 52.9% C and 47.1% O Assume a 100 g sample→

  30. Empirical Formulas • If given the mass of a compound and the mass of the individual elements, then: • Determine how many grams of each element is present • Find the number of moles of each element • Divide each mole number by the smallest mole number • Write the formula • If you come up with a decimal number, multiply the decimal number by a whole number to get a whole number • Then multiply all other ratios by that same whole number

  31. Question • Analysis of 20.0 g of a compound containing only calcium and bromine indicates that 4.00 g of calcium are present. What is the empirical formula of the compound formed?

  32. Molecular Formulas • Defn: The actual number of atoms of each element in one molecule or formula unit of the substance • You calculate a molecular formula from an empirical formula • To find the molecular formula: • find molar mass of empirical formula • Divide molar mass (given) by empirical formula mass to get an “integer” Molar mass of compound = N (an integer) empirical formula mass • Multiply empirical formula by the “integer” • (Molecular Formula) = (Empirical Formula)N… • N distributestoallthesubscripts

  33. Example • The molar mass of a compound is 42 g/mol. Its empirical formula is CH2. What is the molecular formula for this compound?

  34. Example • What is the molecular formula for a compound with an empirical formula of OH and a molar mass of 34 g/mol?

  35. Hydrate Formulas • Defn:a solid which crystallizes with water molecules bonded to the compound in the crystalline lattice. • Name of the hydrous compound: BaCl2 •2H2O → BaCl2(s) + 2 H2O (g) BaCl2 •x H2O → BaCl2(s) + x H2O (g)

  36. To Name a hydrous compound: • Name of the ionic compound + prefix hydrate… (the prefix corresponds to the coefficient in front of the water) • Prefixes: 1-mono 4-tetra 7-hepta 2- di 5-penta 8-octa 3-tri 6-hexa 9-nona 10-deca

  37. Finding a Hydrate Formula • Find the number of moles for each compound • Divide each mol number by the smallest mol number • Write the formula • Name the compound

  38. Example • A hydrate of CuSO4 is heated to drive off the water of crystallization. When 10.0 g of the hydrate is heated, 6.39 grams of solid residue remain. Find the hydrate formula.

  39. The End

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