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Formula and Stoichiometry in Chemical Equations

This article discusses the mathematical relationships involving chemical formulas and equations, atomic mass, formula mass, molecular mass, and the concept of a mole. It also provides examples and conversions related to moles and particles.

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Formula and Stoichiometry in Chemical Equations

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  1. Ch. 8 Formula Stoichiometry Stoichiometry – studies the mathematical relationships involving chemical formulas and equations.

  2. Atomic Mass • The mass of an atom is very very small: • oxygen atom 2.66 x 10-23 grams • hydrogen atom 1.67 x 10-24 grams

  3. Atomic Mass • Although it is possible to measure the actual mass of an atom, it is more convenient to use relative masses. • The relative mass of an atom is a comparison of the atom’s mass to the mass of a carbon-12 atom. • Carbon-12 was chosen as the standard for measuring relative atomic masses.

  4. Atomic Mass Unit • The mass of a carbon-12 atom = exactly 12 u • One atomic mass unit (amu or u) =1/12 the mass of a carbon-12 atom.

  5. Atomic masses for each element are listed on the periodic table.

  6. Some atomic masses: • Hydrogen 1.008 u • Oxygen 16.00 u • Iron 55.85 u What is the atomic mass of: Nitrogen 14.01 u Chlorine 35.45 u Magnesium 24.31 u

  7. Formula Mass • The sum of the atomic masses of all the atoms in a chemical formula. • Formula mass can refer to an element or an ionic or molecular compound. • EX: Formula mass of CO2 is: 12.01 + 2 x 16.00 = 44.01 u

  8. What is the formula mass of Mg(NO2)2? • 1 Mg + 2 N + 4 O • 24.31 u + (2 x 14.01 u) + (4 x 16.00 u) = 116.3 u • Calculate the formula mass of: • 1) NaCl 2) NO2 3) CaCO3 • 4) Ba(NO3)2 5) (NH4)3PO4

  9. Molecular Mass • molecular mass is the formula mass of a molecular substance. It is the mass of a molecule of that substance. • EX: the molecular mass of C6H12O6 = 6 X 12.01 + 12 X 1.008 + 6 X 16.00 • = 180.2 u

  10. The Mole

  11. The Mole: Basic Concepts Measuring Matter • Chemists need a convenient method for counting accurately the number of atoms, molecules, or formula units in a sample of a substance. • As you know, atoms and molecules are extremely small. There are so many of them in even the smallest sample that it’s impossible to actually count them. • That’s why chemists created their own counting unit called the mole.

  12. The mole, commonly abbreviated mol, is the SI base unit used to measure the amount of a substance. • 1 mole (mol) = 6.02 x 1023 particles • “particles” may be atoms, molecules, or formula units The number 6.02 x 1023 is called Avogadro’s number (in honor of Italian Amedeo Avogadro who determined the volume of one mole of a gas.) • If you write out Avogadro’s number, it looks like this: 602 000 000 000 000 000 000 000

  13. A representative particle is any kind of particle such as atoms, molecules, formula units, electrons, or ions. • We will now look at one-mole quantities of three substances, each with a different representative particle.

  14. The representative particle in a mole of water is the water molecule. • 1 mol H2O = 6.02 x 1023 H2O molecules

  15. The representative particle in a mole of copper is the copper atom. • 1 mol Cu = 6.02 x 1023 Cu atoms

  16. The representative particle in a mole of sodium chloride is the formula unit. • 1 mol NaCl =6.02 x 1023 NaClformula units • The term “formula unit”is used to represent one“unit” of an ionic compound.

  17. Mole  Particle Conversions Avagadro's number is used as a conversion factor to convert between # of moles and # of particles (atoms, molecules, ions, etc.)

  18. Converting Moles to Particles • Suppose you want to determine how many particles of sucrose are in 3.50 moles of sucrose. You know that one mole contains 6.02 x 1023 representative particles. • Therefore, you can write a conversion factor, Avogadro’s number, that relates representative particles to moles of a substance.

  19. There are 2.11 x 1024 molecules of sucrose in 3.50 moles of sucrose. EX: How many iron atoms are in 0.025 mol of Fe?

  20. Converting Particles to Moles • Now, suppose you want to find out how many moles are represented by a certain number of representative particles. • You can use the inverse of Avogadro’s number as a conversion factor.

  21. Zinc is used as a corrosion-resistant coating on iron and steel. It is also an essential traceelement in your diet. • Calculate the number of moles that contain 4.50 x 1024 atoms of zinc (Zn).

  22. The Mass of a Mole • Technically, the mole is defined as the number of carbon-12 atoms in exactly 12 g of pure carbon-12. (This has experimentally been shown to be 6.02 x 1023).

  23. Thus, the mass of one mole of carbon-12 atoms is 12 g. What about other elements? (Whether you are considering a single atom or Avogadro’s number of atoms (a mole), the masses of all atoms are established relative to the mass of carbon-12. ) (Play Video) Mass of one mole of: Sulfur 32.06 g Aluminum 26.98 g Sodium 22.99 g H2O 18.02 g O2 32.00 g CO 28.01 g

  24. Molar mass – the mass of one mole of asubstance. The molar mass of a substanceis equal to its formula mass expressed in grams. (units for molar mass are g/mol) (molar mass is sometimes called “gram formula mass” or “gram molecular weight”, etc.)

  25. What is the molar mass of: • iron? • phosphorus? • nitrogen gas (N2)? • sodium chloride? • sulfur dioxide? • Remember: Molar masses are the same as formula mass, except the units are grams; you get molar masses from the periodic table.

  26. Mass(grams)   Moles Conversions • Molar mass is used as a conversion factor to convert between grams (mass) of a substance and # of moles.

  27. Converting Moles to Mass • to convert moles to grams, use the following set up:molar mass • given moles X 1 mol = grams • EX: Calculate the mass of 0.625 moles of calcium.

  28. Converting Mass to Moles • to convert grams to moles, use the following set up:1 mol • given grams X molar mass = moles • EX: Convert 848 g of copper to moles.

  29. Two-Step Problems • converting: # of particles to grams;or grams to # of particles. 6.02 x 1023 X 1 mol 1 mol X molar mass PARTICLES GRAMS (mass) MOLES molar mass X 1 mol 1 mol X 6.02 x 1023

  30. PerkinElmer 2400 Series II CHNS/O Analyzer

  31. Percentage Composition • the percentage composition of a compound is the percentage (by mass) of each element in that compound. • To calculate the percentage of an element in a compound from the compound’s formula: • divide the total atomic masses of all the atoms of that element in the formula BY the formula mass of the formula (then multiply by 100 to get a percentage).

  32. Another type of % composition problem: • Sometimes they just tell you how many grams of each element are in a sample of a compound. • To calculate the percentage of an element in a compound from actual masses: • divide the mass of that element by the total mass of the compound (then multiply by 100 to get a percentage).

  33. Determining Empirical Formulas • In one type of chemical analysis, a compound is decomposed and the masses of the elements that made up the compound are measured. Wethen use the masses of the elements in the compound to calculate the empirical formula.

  34. Steps for Determining Empirical Formula: • Change the mass of each element to moles.(If you are given the percentage of each element, just change the “%” sign to “g”,then convert to moles.) • Divide each answer in step 1 by the smallest answer to obtain the ratio of atoms in the compound. • If step 2 does not give you the simplest WHOLE number ratio, find a number that you can multiply each answer in step 2 by in order to achieve a whole number ratio.

  35. Mass Spectrometer – determines molecular mass

  36. Determining a molecular formula • You can determine the molecular formula from an empirical formula (the molecular mass must be known). • The molecular formula is always a whole number multiple of the empirical formula. • Steps for determining molecular formula: • Determine empirical formula. • Find formula mass of empirical formula. • Divide the given molecular mass by the empirical formula mass. • Multiply each subscript in the empirical formula by the answer in step 3 to give you the molecular formula.

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