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Laws of Chemical
    Combination
        Chem1 SY 10-11
             -kaaferrer
What makes compounds different?
Law of Constant Composition
   1799 Joseph Proust

   a chemical compound contains the same elements
    in exactly the ...
   For example, water always consists of oxygen and
    hydrogen atoms, and it is always 89 percent oxygen
    by mass an...
Does mass change during a chemical
reaction?
Law of Conservation of Mass
   Lavoisier heated a measured amount of mercury to
    form the red oxide of mercury. He mea...
Law of Conservation of Mass
   1744 Antoine Lavoisier
   matter can not be created or destroyed in ordinary
    chemical...
Example

   10 grams of CaCO3 on heating gave 4.4g of CO2
    and 5.6 of CaO. Show that these observations are in
    agr...
Law of Multiple Proportions
   1803 John Dalton

   States that when two elements combine to form
    more than one comp...
Example:
 Carbon monoxide (CO): 12 parts by mass of carbon
  combines with 16 parts by mass of oxygen.

   Carbon dioxid...
   Water has an oxygen-to-hydrogen mass ratio of
    7.9:1.

   Hydrogen peroxide, another compound consisting of
    ox...
Chemical Formula Relationships
How many number of ATOMS?
   Ca(NO)3

   NaNO3

   Ba(IO3 )2
Molecular and Formula Weights
   The sum of all the atomic weights of each atom in its
    chemical formula.



    Formu...
SEATWORK!
Percentage Composition from Formulas
   Percentage by mass contributed by each element in
    the substance.

%element= (...
Stoichiometry

   The quantitative relationships between the
    substances involved in a chemical
    reaction, establis...
Terminologies
     ATOM
      ◦ Smallest particle of an element

     MOLECULE
      ◦ Smallest unit particle of a pure ...
Atomic Mass
   Atomic Mass = used to numerically indicate the mass of
    an atom in its ground state, it is expressed in...
Atomic Mass
          subatomic       charge      Mass (amu)
           particle
           Neutron         None        1....
Atomic Mass vs.
Average Atomic Mass

             Atomic Mass              Average Atomic Mass
  For carbon it is 12 u not...
Average Atomic Mass (calculation)

   Solve for the Average Atomic Mass of the element
    Boron
           Isotope      ...
Relative Atomic Mass &
Atomic Weight

   Have different meaning from atomic mass but
    synonymous with each other, alth...
Molecular Mass
   The average mass of the molecules of a binary
    compound (non metal-non metal)

       Unit: u

    ...
Formula Mass
   The average mass of the molecules of an ionic
    compound (metal-non metal)

       Unit: u

       Ex...
The Mole Concept
     Used to describe the number of particles (atoms,
      molecules, etc.) that make up sample of matt...
Molar Mass

   Mass of 1 mole of substance
   SI unit: g/mol
   Provides a bridge between mass and amount
Conversion



         Use
GRAMS                           FORMULA
        molar   MOLES   Use N
                         ...
Comparison of Atomic and Molar Mass

 Atomic mass                     Molar mass

         Unit:                 Unit:
   ...
Empirical Formula
              and
        Molecular
          Formula
EMPIRICAL FORMULA


   Shows the relative number of atom of each
    element in the compound
Problem 1
     A compound is found out to contain
 20.0% carbon, 2.2% hydrogen, and 77.8%
 chlorine. Determine EF of the c...
MOLECULAR FORMULA


   Shows the actual number of atom of each
    element in the compound
Relationship between EF & MF

              MF = (EF)n
where n is an integer
therefore,
               n = MMF
           ...
Problem 2
    A compound with the empirical formula
 C2H5 has a molar mass of 58.12 g/mol.
 Find the molecular formula of ...
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 Mole

   One mole is the amount of any substance that
    contains the same number of units as the number of
    atom...
MOLAR Mass
   The mass in grams of one (1) mole of a substance.
   g/mol, grams per mole




   The molar mass of any s...
Avogadro
           ’s             Formula
atoms                               gram
                   mole    weight
    ...
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Laws of Chemical Combination + Stoichiometry Slide 1 Laws of Chemical Combination + Stoichiometry Slide 2 Laws of Chemical Combination + Stoichiometry Slide 3 Laws of Chemical Combination + Stoichiometry Slide 4 Laws of Chemical Combination + Stoichiometry Slide 5 Laws of Chemical Combination + Stoichiometry Slide 6 Laws of Chemical Combination + Stoichiometry Slide 7 Laws of Chemical Combination + Stoichiometry Slide 8 Laws of Chemical Combination + Stoichiometry Slide 9 Laws of Chemical Combination + Stoichiometry Slide 10 Laws of Chemical Combination + Stoichiometry Slide 11 Laws of Chemical Combination + Stoichiometry Slide 12 Laws of Chemical Combination + Stoichiometry Slide 13 Laws of Chemical Combination + Stoichiometry Slide 14 Laws of Chemical Combination + Stoichiometry Slide 15 Laws of Chemical Combination + Stoichiometry Slide 16 Laws of Chemical Combination + Stoichiometry Slide 17 Laws of Chemical Combination + Stoichiometry Slide 18 Laws of Chemical Combination + Stoichiometry Slide 19 Laws of Chemical Combination + Stoichiometry Slide 20 Laws of Chemical Combination + Stoichiometry Slide 21 Laws of Chemical Combination + Stoichiometry Slide 22 Laws of Chemical Combination + Stoichiometry Slide 23 Laws of Chemical Combination + Stoichiometry Slide 24 Laws of Chemical Combination + Stoichiometry Slide 25 Laws of Chemical Combination + Stoichiometry Slide 26 Laws of Chemical Combination + Stoichiometry Slide 27 Laws of Chemical Combination + Stoichiometry Slide 28 Laws of Chemical Combination + Stoichiometry Slide 29 Laws of Chemical Combination + Stoichiometry Slide 30 Laws of Chemical Combination + Stoichiometry Slide 31 Laws of Chemical Combination + Stoichiometry Slide 32 Laws of Chemical Combination + Stoichiometry Slide 33 Laws of Chemical Combination + Stoichiometry Slide 34 Laws of Chemical Combination + Stoichiometry Slide 35 Laws of Chemical Combination + Stoichiometry Slide 36 Laws of Chemical Combination + Stoichiometry Slide 37 Laws of Chemical Combination + Stoichiometry Slide 38 Laws of Chemical Combination + Stoichiometry Slide 39
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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 chemical reaction?
  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 Use GRAMS 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)n where n is an integer therefore, n = MMF MEF where 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 Formula atoms gram mole weight number
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