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Chemistry 30 – Review of Basic Chemistry 20

Chemistry 30 – Review of Basic Chemistry 20. Polyatomic Molecular Elements:. Naming Molecular Compounds:. Combining elements of the periodic table that come only from the nonmetals (right side of the “staircase” only) forms a molecular compound. Common molecular compounds include:.

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Chemistry 30 – Review of Basic Chemistry 20

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  1. Chemistry 30 – Review of Basic Chemistry 20

  2. Polyatomic Molecular Elements:

  3. Naming Molecular Compounds: Combining elements of the periodic table that come only from the nonmetals (right side of the “staircase” only) forms a molecular compound. Common molecular compounds include:

  4. When naming molecular compounds: • write the first name as given on the periodic table of elements. • write the last name using an “ide” ending. • place the appropriate prefix in front the first and last name to describe the number of atoms there are of each element. • where the first element has only one atom, “mono” is not necessary. Example: P4O3(g) = 4 atoms tetraphosphorus trioxide 3 atoms

  5. Naming Ionic Compounds: • Combining elements of the periodic table that come from the metalsandnonmetals (left and right side of the “staircase” only) forms an ionic compound. • When naming ionic compounds: • write the first name as given on the periodic table of elements. • write the last name using an “ide” ending. • use no prefixes. Example: CaCl2 - calcium chloride

  6. When writing simple ionic formulas: • put down the metallic element first. • put down the nonmetallic element last. • cross the elements’ ionic charge to become the subscript for each other element. • numerically simplify the subscripts. Example: magnesium phosphide - Mg2+ and P3– join to produce MgP Use the charge of one element to be the subscript for the other element – Mg3P2 Example: calcium oxide - Ca2+ and O2– join to produce CaO Use the charge of one element to be the subscript for the other element – Ca2O2 Now simplify – CaO all ionic compounds are solids at room temperature.

  7. When writing ionic formulas involving complex ions: • use the same format as above but whenever a complex ion is named, use brackets to keep that complex ion as a group. Example: sodium sulfate - Na+ and SO42– Put the two together grouping the complex ion: Na+ (SO42–) Now cross the charges: Na2(SO4)1 Since 1’s are not necessary: Na2SO4 Example: calcium nitrate - Ca2+ and NO3– Ca2+(NO3–) Ca(NO3)2 Example: sodium hydroxide - Na+ and OH– Na+(OH–) NaOH

  8. When writing ionic formulas involving elements with more than one charge: • use the first ion listed as the most common. For example, Cu2+ is more common than Cu+, so Cu2+ would be used if no choice is given. When naming these compounds containing elements with more than one charge: • use Roman numerals to indicate the charge of the ion used. • Example: CuCl is copper (I) chloride • Example: CuCl2 is copper (II) chloride

  9. Hydrated Compounds • When writing hydrated compounds, follow all ionic rules described above. Then use a dot along with the number of water molecules required. • When naming hydrated compounds, follow all ionic rules described above. Then use a prefix in front of the word “hydrate”. Example: CuSO4 6H2O is: copper (II) sulfate hexahydrate Example: aluminum chloride trihydrate is: AlCl3 3H2O

  10. Clues for a Chemical Reaction • Formation of a precipitate • Formation of a gas • Colour change • Energy change

  11. Types of Chemical Reactions Simple Composition element + element  compound 2 Na(s) + Br2(g)  Simple Decomposition compound  element + element + element 2CaCO3(s)  Single Replacement element + compound  element + compound Mg(s) + 2 NaOH(aq)  Double Replacement compound + compound  compound + compound 3 HCl(aq) + Al(OH)3(aq)  Hydrocarbon Combustion hydrocarbon + oxygen  carbon dioxide + water vapour C3H8(g) + O2(g)  2 NaBr(s) 2Ca(s) + 2C(s) + 3O2(g) 2 Na(s) + Mg(OH)2(aq) AlCl3(s) + 3 HOH(l) 3 CO2(g) + 4 H2O(g)

  12. Writing Dissociation Equations Compounds that dissolve in water may produce ions. These solutions are called electrolytes. Some compounds may dissolve in water but form no ions. These solutions are called nonelectrolytes. When electrolytes are formed, dissociation equations can be shown. Examples: NaOH(aq)  Na+(aq) + OH–(aq) Al2(SO4)3(aq)  2 Al3+(aq) + 3 SO42–(aq) Use mole ratios to determine solution or ion concentrations. Examples: NaOH(aq)  Na+(aq) + OH–(aq) 2.0 mol/L ? ? 2.0 mol/L 2.0 mol/L Al2(SO4)3(aq)  2 Al3+(aq) + 3 SO42–(aq) ? 3.0 mol/L ? 1.5 mol/L 4.5 mol/L

  13. Writing Nonionic, Total Ionic and Net Ionic Equations Example: A silver nitrate solution reacts with a solution of barium chloride. AgNO3(aq) + BaCl2(aq)  Ba(NO3)2(aq) + AgCl(s) (unbalanced) Nonionic Equation: (regular balanced equation) 2 AgNO3(aq) + BaCl2(aq)  Ba(NO3)2(aq) + 2 AgCl(s) Total Ionic Equation: (list dissociations for electrolytes only) 2 Ag+(aq) + 2 NO3–(aq) + Ba2+(aq) + 2 Cl–(aq)  Ba2+(aq) + 2 NO3–(aq) + 2 AgCl(s) (do not write dissociations for solids, liquids or gases) Net Ionic Equation: (list only what reacts or changes) 2 Ag+(aq) + 2 Cl–(aq)  2 AgCl(s) or, simplified: Ag+(aq) + Cl–(aq)  AgCl(s)

  14. Significant Digits • All numbers listed are significant except zeros before or after a decimal that must be used as placeholders. Example: 100.0010 - 7 significant digits 0.001010 - 4 significant digits

  15. Significant DigitsContinued… • Multiplication or Division Rules: • Count the number of digits in each number being multiplied or divided. • Perform the multiplication or division. • Round off to the least number of digits found in each of the individual numbers being multiplied.

  16. Significant DigitsContinued… 0.094 (answer to 2 sig. dig) Example: 2.34 x 3.342 x 0.012 = 0.09384336 = 3 4 2 Example: 3.54 x 120.4 x 0.10 = 42.6216 = 3 4 2 Example: 35.127 x 225.5 x 2.75 = 21783.13088 = 5 4 3 Example: 350.55  12 = 29.2125 = 2 Example: 12  350 = 0.034285714 = 2 3 43 2.18 x 104 29 0.034

  17. Addition or Subtraction Rules: Count the number of digits following the decimal in each number being added or subtracted. Perform the addition or subtraction. Round off to the least number of digits found following the decimal. Example: 2.34 (2) + 2.8 (1)5.14 =5.1 Example: 3.54 (2) – 1.134 (3) 2.406 =2.41 Note: Always save all number in the calculator and round off only for your final answer.

  18. Stoichiometry • Determine the balanced chemical equation. • Determine information given. • Determine what it is you are solving for. • Determine the number of moles of what is given. • Use a mole ratio to determine the number of moles of the unknown. • Solve for the answer. Example: If 200 mL of 0.100 mol/L silver nitrate solution reacts with a piece of copper, determine the mass of metal reacted. 2 AgNO3 (aq) + Cu(s) 2 Ag(s) + Cu(NO3)2(aq) v = 0.200 L m = ? C = 0.100 mol/L nAgNO3 = Cv nAgNO3= (0.200 mol/L)(0.100 L) nAgNO3 = nCu = 0.0200 mol x ½ = mCu = nM mCu = (0.0100 mol)(63.55 g/mol) mCu = 0.0200 mol 0.0100 mol 0.636 g

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