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Laws of Chemical Combination + Stoichiometry

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Laws of Chemical Combination + Stoichiometry

  1. 1. Laws of Chemical Combination Chem1 SY 10-11 -kaaferrer
  2. 2. What makes compounds different?
  3. 3. Law of Constant Composition 1799 Joseph Proust a chemical compound contains the same elements in exactly the same proportions (ratios) by mass regardless of the size of the sample or source of the compound
  4. 4.  For example, water always consists of oxygen and hydrogen atoms, and it is always 89 percent oxygen by mass and 11 percent hydrogen by mass
  5. 5. Does mass change during a chemicalreaction?
  6. 6. Law of Conservation of Mass Lavoisier heated a measured amount of mercury to form the red oxide of mercury. He measured the amount of oxygen removed from the jar and the amount of red oxide formed. When the reaction was reversed, he found the original amounts of mercury and oxygen.
  7. 7. Law of Conservation of Mass 1744 Antoine Lavoisier matter can not be created or destroyed in ordinary chemical or physical changes. the mass of the reactants (starting materials) equals the mass of the products 2Mg (s) + O2 (g) → 2MgO (s) 48.6 g 32.0 g 80.6 g
  8. 8. Example 10 grams of CaCO3 on heating gave 4.4g of CO2 and 5.6 of CaO. Show that these observations are in agreement with the law of conservation
  9. 9. Law of Multiple Proportions 1803 John Dalton States that when two elements combine to form more than one compound, the masses of one element which combine with a fixed mass of the other element are in ratios of small whole numbers
  10. 10. Example: Carbon monoxide (CO): 12 parts by mass of carbon combines with 16 parts by mass of oxygen. Carbon dioxide (CO2): 12 parts by mass of carbon combines with 32 parts by mass of oxygen. Ratio of the masses of oxygen that combines with a fixed mass of carbon (12 parts) 16: 32 or 1: 2
  11. 11.  Water has an oxygen-to-hydrogen mass ratio of 7.9:1. Hydrogen peroxide, another compound consisting of oxygen and hydrogen, has an oxygen-to-hydrogen mass ratio of 15.8:1. Ratio of the masses of oxygen that combines with a fixed mass of hydrogen is 7.9: 15.8 or 1: 2
  12. 12. Chemical Formula Relationships
  13. 13. How many number of ATOMS? Ca(NO)3 NaNO3 Ba(IO3 )2
  14. 14. Molecular and Formula Weights The sum of all the atomic weights of each atom in its chemical formula. Formula mass = ∑ atomic masses in the formula unit
  15. 15. SEATWORK!
  16. 16. Percentage Composition from Formulas Percentage by mass contributed by each element in the substance.%element= (# of atoms of that element)(atomic mass of that element) ------------------------------------------------------------------------ formula mass of the compound
  17. 17. Stoichiometry The quantitative relationships between the substances involved in a chemical reaction, established by the equation for the reaction
  18. 18. Terminologies  ATOM ◦ Smallest particle of an element  MOLECULE ◦ Smallest unit particle of a pure substance  ION ◦ An atom or group of bonded atoms with electrical charge because of an excess or deficiency of electrons  ELEMENT ◦ Pure substance; CANNOT be broken down into 2 or more pure substances by chemical means  COMPOUND ◦ Pure substance; CAN be broken down into 2 or more pure substances by chemical means
  19. 19. Atomic Mass Atomic Mass = used to numerically indicate the mass of an atom in its ground state, it is expressed in the non SI unit of u u = refers to unified atomic mass unit (formerly known as atomic mass unit or amu) 1 amu = 1/12 the mass of carbon-12 atom, therefore the mass of C-12 atom is made EQUAL to 12 amu Carbon-12 atom is an isotope of carbon 1 amu = 1.66 x 10-24 g
  20. 20. Atomic Mass subatomic charge Mass (amu) particle Neutron None 1.0087 ≈ 1 Proton Positive 1.0073 ≈ 1 Electron Negative 5.486 x 10-4 ≈ 0  Mass of e = 1/1800 of mass of p and n so it is negligible making the equation Atomic mass of 12C = mass of p + mass of n
  21. 21. Atomic Mass vs.Average Atomic Mass Atomic Mass Average Atomic Mass For carbon it is 12 u not 12.01 u *Used to relate the fact that the numerical value assigned to each element in the periodic table reflects the average abundances of the atoms that compose a naturally occurring element *Related to isotopes *For carbon it is 12.01 u *Chemists often will use the term “atomic mass” when they are actually referring to average atomic mass of an atom.
  22. 22. Average Atomic Mass (calculation) Solve for the Average Atomic Mass of the element Boron Isotope Mass (u) Percent Abundance 11B 11.009305 80.1 10B 10.012937 19.9 Average atomic mass = ∑ (mass x percent abundance) Where ∑ means “sum”
  23. 23. Relative Atomic Mass &Atomic Weight Have different meaning from atomic mass but synonymous with each other, although of different historical origins Relative Atomic Mass  “Ratio” of the average mass of the atom to the unified atomic mass; dimensionless Atomic Weight  The name Dalton used in the early 19th century to numerically describe the weight of atoms relative to each other
  24. 24. Molecular Mass The average mass of the molecules of a binary compound (non metal-non metal)  Unit: u  Ex.: The molecular mass of carbon dioxide gas, CO2 is 28.01 u.
  25. 25. Formula Mass The average mass of the molecules of an ionic compound (metal-non metal)  Unit: u  Ex.: The formula mass of barium chloride, BaCl2 is 208.2 u.
  26. 26. The Mole Concept  Used to describe the number of particles (atoms, molecules, etc.) that make up sample of matter “One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12.” 1 mol of any substance = 6.02 x 1023 units of that substance, where 6.02 x 1023 Avogadro’s Number, N
  27. 27. Molar Mass Mass of 1 mole of substance SI unit: g/mol Provides a bridge between mass and amount
  28. 28. Conversion UseGRAMS FORMULA molar MOLES Use N UNITS mass
  29. 29. Comparison of Atomic and Molar Mass Atomic mass Molar mass Unit: Unit: amu; u g/mol Atomic Numerically Macroscopic level equivalent level Total mass Mass of 1 of p+ and no mole
  30. 30. Empirical Formula and Molecular Formula
  31. 31. EMPIRICAL FORMULA Shows the relative number of atom of each element in the compound
  32. 32. Problem 1 A compound is found out to contain 20.0% carbon, 2.2% hydrogen, and 77.8% chlorine. Determine EF of the compound.
  33. 33. MOLECULAR FORMULA Shows the actual number of atom of each element in the compound
  34. 34. Relationship between EF & MF MF = (EF)nwhere n is an integertherefore, n = MMF MEFwhere M is molar mass
  35. 35. Problem 2 A compound with the empirical formula C2H5 has a molar mass of 58.12 g/mol. Find the molecular formula of the compound.
  36. 36. Problem 3 The molar mass of a compound that is 54.6% carbon, 9.0% hydrogen and 36.4% oxygen is 88 g/mol. Find MF of the compound.
  37. 37. The Mole One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12. 6.02 x 1023  Avogadro’s Number
  38. 38. MOLAR Mass The mass in grams of one (1) mole of a substance. g/mol, grams per mole The molar mass of any substance in grams per mole is always numerically equal to the atomic, formula, or molecular mass of the substance in amu.
  39. 39. Avogadro ’s Formulaatoms gram mole weight number

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