Stoichiometry Part I

  • 1,737 views
Uploaded on

Atomic Mass up to Molecular Formula

Atomic Mass up to Molecular Formula

More in: Education , Technology
  • Full Name Full Name Comment goes here.
    Are you sure you want to
    Your message goes here
  • the content might look disorganized (idk why) however, the downloaded version is alright. :)
    Are you sure you want to
    Your message goes here
No Downloads

Views

Total Views
1,737
On Slideshare
0
From Embeds
0
Number of Embeds
0

Actions

Shares
Downloads
76
Comments
1
Likes
4

Embeds 0

No embeds

Report content

Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
    No notes for slide

Transcript

  • 1. Ariane B. Rosos Chemistry 1 SY 2009 - 2010
  • 2.
    • The quantitative relationships between the substances involved in a chemical reaction, established by the equation for the reaction
  • 3.
    • 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
  • 4.
    • 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
  • 5.
    • Mass of e = 1/1800 of mass of p and n so it is negligible making the equation
    • Atomic mass of 12 C = mass of p + mass of n
    subatomic particle charge Mass (amu) Neutron None 1.0087 ≈ 1 Proton Positive 1.0073 ≈ 1 Electron Negative 5.486 x 10 -4 ≈ 0
  • 6. 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.
  • 7.
    • Solve for the Average Atomic Mass of the element Boron
    • Average atomic mass
    • = ∑ (mass x percent abundance)
    • Where ∑ means “sum”
    Isotope Mass (u) Percent Abundance 11 B 11.009305 80.1 10 B 10.012937 19.9
  • 8.
    • 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 19 th century to numerically describe the weight of atoms relative to each other
  • 9.
    • The average mass of the molecules of a binary compound (non metal-non metal)
      • Unit: u
      • Ex.: The molecular mass of carbon dioxide gas, CO 2 is 28.01 u.
  • 10.
    • The average mass of the molecules of an ionic compound (metal-non metal)
      • Unit: u
      • Ex.: The formula mass of barium chloride, BaCl 2 is 208.2 u.
  • 11.
    • 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 10 23 units of that substance,
    • where 6.02 x 10 23 Avogadro’s Number, N
  • 12.
    • Mass of 1 mole of substance
    • SI unit: g/mol
    • Provides a bridge between mass and amount
  • 13. GRAMS MOLES FORMULA UNITS Use molar mass Use N
  • 14. Atomic mass Molar mass Unit: g/mol Unit: amu; u Atomic level Macroscopic level Numerically equivalent Comparison of Atomic and Molar Mass Total mass of p + and n o Mass of 1 mole
  • 15.
    • Hydrogen
    • Atomic mass: 1 atom H = 1.008 u
    • 1 atom =
    • Molar mass: 1 mole H =
  • 16.
    • Avogadro’s number, N = 6.0221415 × 10 23 particles
    • 1 u = 1.660538782 × 10 −24 grams
    • 0.999999913 ≈ 1
  • 17.  
  • 18.
    • Shows the relative number of atom of each element in the compound
  • 19.
    • A compound is found out to contain 20.0% carbon, 2.2% hydrogen, and 77.8% chlorine. Determine EF of the compound.
  • 20.
    • Shows the actual number of atom of each element in the compound
  • 21.
    • MF = (EF)n
    • where n is an integer
    • therefore ,
    • n = M MF
    • M EF
    • where M is molar mass
  • 22.
    • A compound with the empirical formula C 2 H 5 has a molar mass of 58.12 g/mol. Find the molecular formula of the compound.
  • 23.
    • 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.
  • 24.
    • Thursday, 16 Sept 2009
    • Bring:
      • Scientific calculator
      • Black/blue ball pen
      • Intermediate paper
  • 25.
    • Stoichiometry I - Part I and Part II
    • (up to page 8)
    • Intermediate paper; show complete solution and express final answers in correct significant figures.
    • Consultation (optional) - Tuesday, 15 September, 12nn-2:30pm or through email: [email_address]
    • Deadline: Wednesday, 16 September 2009
    • 12 noon, Chemistry Unit
    • Answer key: available 2pm (Wed) @ the bulletin board in front of chemistry unit 