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Chapter 5

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Chapter 5

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  1. Chapter 5 Elaine Teto

  2. The Mole • the formula mass in grams of a substance contains one mole of particles • A mole is also equivalent to Avogadro’s number (6.022 x 1023­.) • The formula would be: mole = n = mass (gr) / molar mass (gr/mole)

  3. Returning back to what was previously stated a mole of anything has Avogadro’s number of objects in it • For example, a mole of water (H2O) has 6.022 x 1023 water molecules. • a mole of 12C atoms will weigh 12 grams, a mole of 197Au atoms will weigh 197 grams.

  4. Atomic mass, formula weight and molar mass • atomic mass unit (amu) is defined by setting the mass of one C atom equal to 12 amu • We can determine the amu experimentally to be 1.66 x 10-24 grams. • These weights can be derived from the periodic table

  5. the formula weight, which is the sum of the atomic weight of each atom in a chemical formula. • For example, the formula weight for NaOH is (22.99 + 15.99 + 1.0079) 39.988 amu. • When chemical formulas equal the molecular formula than the formula weight is also known as the molecular weight.

  6. the molar mass, the mass in grams of 1 mole of a substance, can be determined. • The molar mass is numerically equal to the formula weight. For example one mole of water (FW = 18 amu) has a molar mass of 18 grams.

  7. Stoichiometry • Stoichiometry is the math behind chemistry. • Given enough information, one can use stoichiometry to calculate masses, moles, and percents within a chemical equation. • when considering balancing equations one must remember the Law of Conservation of Mass, which states that matter can neither be created nor destroyed

  8. Balancing Equations • Example 1: Al + Fe3O4  Al2O3 + Fe • When balancing equations it is important to take one atoms at a time and work with each individually. Trying to balance an entire equation at once can be difficult. • 8 Al + 3 Fe3O4  4 Al2O3 + 9 Fe • Example 2: CH4 + 2O2  CO2 + H2O + Heat • CH4 +2O2  CO2 + 2H2O + Heat

  9. Empirical and Molecular formulas • The molecular formula is the form of the term as it would appear in a chemical equation. The empirical formula and the molecular formula can be the same, or the molecular formula can be any positive integer multiple of the empirical formula. • Examples of empirical formulas: AgBr, Na2S, C6H10O5. • Examples of molecular formulas: P2, C2O4, C6H14S2, H2, C3H9

  10. One can calculate the empirical formula from the masses or percentage composition of any compound. • Example: Calculate the empirical formula for a compound that has 43.7 g P (phosphorus) and 56.3 grams of oxygen. • First we convert to moles: • Next we divide the moles to try to get an even ratio. • When we divide, we did not get whole numbers so we must multiply by two (2). • The answer=P2O5

  11. After one has calculated the empirical formula one can calculate the molecular formula. By dividing the molecular mass of the compound by the mass of the empirical formula, one would attain the molecular formula

  12. Example: The empirical formula of a compound is HCN, 2.016 grams of hydrogen are necessary to make the compound, what is the molecular formula? • In the empirical formula hydrogen weighs 1.008 grams. Dividing 2.016 by 1.008 we see that the amount of hydrogen needed is twice as much. Therefore the empirical formula needs to be increased by a factor of two (2). The answer is: H2C2N2.

  13. Concentrations of solutions • The concentration of a solution is typically given in molarity. Molarity is defined as the number of moles of solute (what is actually dissolved in the solution) divided by the liters of solution (the total volume of what is dissolved and what it has been dissolved in).

  14. Molarity = moles of solute liters of solution • Example: If 5.00 grams of NaOH are dissolved in 5000 mL of water, what is the molarity of the solution? • the molarity (M) of the solution is 0.025 mol/L.

  15. Molality is another common measurement of concentration. Molality is defined as moles of solute divided by kilograms of solvent (the substance in which it is dissolved, like water). • It is possible to convert between molarity and molality. The only information needed is density.

  16. Chemical Reactions • one of the best solvents is water, because it is a polar molecule • Electrolytes are substances whose aqueous solutions conduct electricity • When an ionic substance is broken apart into the solvated cations and anions, the process is properly called dissociation. However, when the ions are formed by the reaction of the substance with water, and didn't exist in the original substance, then the process is properly called ionization.

  17. Some examples of substances which are classified as strong electrolytes are: • * strong inorganic acids (the common ones are: hydrochloric acid, hydrobromic acid, hydroiodic acid, sulfuric acid, nitric acid, chloric acid, and perchloric acid • * strong inorganic bases, e.g. NaOH. • Dissociation reaction: NaCl + H2O ---> Na+ (aq) + Cl- (aq) • Ionization Reaction: HCl(g) + H2O(l) ---> H3O+ (aq) + Cl- (aq)

  18. Some examples substances which behave as weak electrolytes are: • * Weak inorganic acids (inorganic acids which are not listed above), e.g. carbonic acid, nitrous acid, hypochlorous acid, etc. • * Organic acids (carboxylic acids, e.g. acetic acid) • * Organic amines, e.g. methylamine.

  19. equations for chemical reactions • Molarity equation (Molarity) x (Volume [in liters]) = moles of solute • Dilution equation MCVC = MDVD

  20. Oxidation states • Oxidation is the loss of electrons and reduction is the gain of electrons • The method used for keeping track of the electrons is the oxidation number. • The only time that an oxidation number is a real charge is for monatomic ions, like Cl- (oxidation number -1) or aluminum ion, Al3+ (oxidation number +3). In hydrogen chloride, the oxidation numbers of +1 for the hydrogen and -1 for the chlorine are clearly not real; HCl is a covalent compound which doesn't contain ions.

  21. Conclusion • Chapter five contained many significant topics in the study of chemistry. • It is important to remember that, starting from the simplest form of a concept, such as the mole, and than working toward a more elaborate concept, such as stoichiometry, is the best way to master the topics.