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CHAPTER 2

CHAPTER 2. CHEMICAL FORMULAS & COMPOSITION STOICHIOMETRY. Chemical Formulas. show the ratio of the elements present in the molecule or compound He, Au, Na - monatomic O 2 , H 2 , Cl 2 - diatomic O 3 , P 4 , S 8 - more complex elements

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CHAPTER 2

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  1. CHAPTER 2 CHEMICAL FORMULAS & COMPOSITION STOICHIOMETRY

  2. Chemical Formulas • show the ratio of the elements present in the molecule or compound • He, Au, Na - monatomic • O2, H2, Cl2 - diatomic • O3, P4, S8 - more complex elements • H2O, C12H22O11 – compounds ~ contains 2 or more elements • Allotropes ~ different forms of the same element in the same physical state. Ex. Both O2 and O3 are gases

  3. Compound • HCl • H2O • NH3 • C4H10 1 Molecule Contains • 1 H atom & 1 Cl atom ratio of elements: 1/1 • 2 H atoms & 1 O atom ratio of elements: 2/1 • 1 N atom & 3 H atoms ratio of elements: 1/3 • 4 C atoms & 10 H atoms ratio of elements: 4/10 = 2/5

  4. Ions & Ionic Compounds • ions are atoms or groups of atoms with an electrical charge • two basic types of ions • positive ions or cations (+) • one or more electrons less than neutral, ex. Na+ • negative ions or anions (-) • one or more electrons more than neutral, ex. Cl- • cations + anions must give neutral charge • KOH potassium hydroxide (+1 & -1) • CaSO4 calcium sulfate (+2 & -2) • Sr3N2 strontium nitride ((2x3)&(-3x2)) • Al(OH)3 aluminum hydroxide (+3 & (-1x3))

  5. Ions & Ionic Compounds • Sodium chloride - table salt is an ionic compound • When a soluble ionic compound is dissolved in water, it dissociates (breaks apart) into its ions. Ex. Salt in water, Na+ and Cl-

  6. Atomic Weights • Scientists made a scale for comparing the masses of all elements. The units are arbitrary ~ called atomic mass units (amu) • How? Weighted average of the masses of the constituent isotopes, its the lower number on periodic chart • 1 amu = 1/12 mass of Carbon-12 • On the periodic table, • atomic weight of H = 1.0079 amu • atomic weight of Ca = 40.078 amu • Ca atom has 40 times more mass than the H atom

  7. The Mole • Atoms are very small, difficult to weigh and count, so how can we measure them accurately? • Tha MOLE!!! • strictly a convenience unit • amount that is large enough to see and handle in lab • mole = number of things • dozen = 12 things • mole = 6.022 x 1023 things • Avogadro’s number = 6.022 x 1023

  8. The atomic weight (amu) = mass of 1 mol (g), also called molar mass - mass in grams equal to the atomic weight • Units of atomic weight are amu/atoms or g/mol • He exists as atoms: • 4.0026 g of He atoms = 1 mol He = 6.022 x 1023 atoms He • H2 exists as molecules: • 2.0158 g of H2 molecules= 1 mol H2 = 6.022 x 1023 molecules H2 = 2 (6.022 x 1023)atoms of H = 1.204 x 1024 atoms of H

  9. Ex. 1) Calculate the mass of a single K atom in grams to 4 significant figures. • Ex. 2) Calculate the number of atoms in one-millionth of a gram of K to 4 significant figures. • Ex. 3) What element do you have if 3.00 moles of it weighs 80.9 g? • Ex. 4) How many moles of Mercury are contained in 9.84mL of Hg? How many atoms? Specific gravity of Hg = 13.5939

  10. Formula and molecular weights • Formula weight (for ionic compounds) formula units, molecular weight (for covalent compounds) molecules – sum of masses of elements in a compound • Ex. 5) The molar mass of calcium nitrate is: • Ex. 6) Calculate the formula weight of hydrated Al2(SO4)3.18H2O. Hydrates have water attached. • Ex. 7) How many a) moles, b) molecules, c) oxygen atoms are contained in 70.0 g of ozone, O3?

  11. Ex. 8) Calculate the number of Na atoms in 36.50 g of Na2CO3. • Ex. 9) What mass of ammonium phosphate, (NH4)3PO4, would contain 15.0 g of N? B) If you start with 15.0 g of H, how much (NH4)3PO4 would you make? • Ex. 10) Calculate the number of mmol in 0.23 g of oxalic acid, (COOH)2.

  12. Percent Composition • mass of each element divided by the mass of the whole compound x 100% • Ex. 11) Find the % by mass of the elements in hydrated FeSO4 .7H2O. Assume you have 1 mole of the compound. • All samples of hydrated iron(II) sulfate have this composition b/c of Law of Definite Proportions which states that different samples of a pure compound contains the same elements in the same proportions.

  13. empirical & molecular formulas • Many times in the lab, a chemist will synthesize a compound. To help prove what was made the compound is sent for elemental analysis. From this the simplest formula is found. • empirical formula - simplest molecular formula, shows ratios of elements but not actual numbers of elements • molecular formula - actual numbers of atoms of each element in the compound • determine empirical & molecular formulas of a compound from percent composition • percent composition is determined experimentally

  14. empirical & molecular formulas • Ex. of molecular and empirical formulas • Molecular Empirical C6H6 CH P4 O10 P2O5 SO2SO2 (same) • Ex. 12) A compound is found to contain 85.63% C and 14.37% H by mass, what is it’s empirical formula? In another experiment its molar mass is found to be 56.1 g/mol. What is its molecular formula?

  15. Purity of Samples • The percent purity of a sample of a substance is always represented as % purity = mass of pure substance x 100% mass of sample ~ mass of sample includes impurities (works just like percentages) • Ex. 13) A bottle of sodium phosphate, Na3PO4, is 92.3% pure Na3PO4. What are the masses of Na3PO4 and impurities in 250. g of this sample of Na3PO4?

  16. In 1986, Bednorz and Muller succeeded in making the first of a series of chemical compounds that were superconducting at relatively high temperatures. This first compound was La2CuO4 which superconducts at 35K. In their initial experiments, Bednorz and Muller made only a few mg of this material. How many La atoms are present in 3.56 mg of La2CuO4?

  17. Within a year after Bednorz and Muller’s initial discovery of high temperature superconductors, Wu and Chu had discovered a new compound, YBa2Cu3O7, that began to superconduct at 100 K. If we wished to make 1.00 pound of YBa2Cu3O7, how many grams of yttrium must we buy?

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