1. Unit One Notes

2. Nomenclature Review ● Ionic ● metal/nonmetal ● Names – ide ending – roman numerals ● charges ● Magnesium fluoride ● copper(II) hydroxide ● FeO ● Ca 3 P 2 ● Ca 3 (PO 4 ) 2

3. Nomenclature Review ● Covalent – 2 nonmetals – Prefixes – ide endings ● Carbon dioxide ● Phosphorus trichloride ● P 2 O 5 ● OF 2

4. Nomenclature Review ● Acids ● Binary – Hydro ___ ic acid ● Oxyacids – ate = ___ic acid – ite = ___ous acid ● Hydrochloric acid ● Hydrosulfuric acid ● Sulfurous acid ● Sulfuric acid ● HF ● H 3 PO 4

5. Metric System ● M = mega = 1x10 6 ● k = kilo = 1000 ● c = centi = 1/100 ● m = milli = 1/1000 ● μ = micro = 1x10 -6 ● n = nano = 1x10 -9 ● 1Mg = 1x10 6 g ● 1kg = 1000g ● 1g = 100cg ● 1g = 1000mg ● 1g = 1x10 6 μg ● 1g = 1x10 9 ng

6. Significant Figures ● Used to determine precision of data – Last digit is estimated ● Example: 2.3g or 2.304g – Which value is more precise? – Which number do you use? Why? ● When data is given sig figs determine how answer is rounded – Most answers will be rounded to 3 sig figs

7. Sig Fig Rules ● Multiply/divide: answer has least number of sig figs ● Add/subtract: answer has fewest number of decimal places ● What is significant? – All non-zeroes – Zeroes between non-zeroes – Zeroes in front never significant – Trailing zeroes are significant if after a decimal

8. Sig Fig Practice

9. Scientific Notation ● Can show sig figs – 1.00 x 10 3 = 1000 – How many sig figs in each number? ● Use EE or EXP button on calculator instead of parentheses and ^

10. Problem Solving ● Read problem – Identify knowns and unknowns ● Plan – Decide which equation to use – Do you need data from other sources? ● Compute – Manipulate equation so the unknown variable is on one side – UNITS, UNITS, UNITS!!! ● Evaluate – Does the answer make sense? – Are the units correct?

11. Precision vs. Accuracy ● Precision: how close the measurements are to each other ● Accuracy: how close the measurements are to a known or accepted value ● Percent Error = ● Several trials for precision ● Use average to determine accuracy (% error)

12. Isotope Notation

13. Mass Number vs. Average Atomic Mass ● Mass Number: sum of p and n in an isotope ● Average Atomic Mass: average of all naturally occurring isotopes – Mg has 3 isotopes: ● 78.99% is 23.985amu ● 10.00% is 24.986amu ● 11.01% is 25.982amu – Calculate average atomic mass

14. Calculating with Moles ● 1mole = molar mass ● 1mole = 6.02x10 23 particles ● How many moles in 18.5g CaCl 2 ? ● How many grams in 5.6x10 24 molecules CH 4 ?

15. Empirical Formulas ● Determined by lab data – Can determine molecular formula if you have the mass ● Percent to mass ● Mass to mole ● Divide by small ● Multiply till whole ● All ionic compounds are empirical ● Covalent compounds may or may not be empirical

16. Empirical Formulas ● Determined by lab data – Can determine molecular formula if you have the mass ● Percent to mass ● Mass to mole ● Divide by small ● Multiply till whole ● All ionic compounds are empirical ● Covalent compounds may or may not be empirical

17. Empirical Examples ● Empirical or Molecular? – H 2 O – H 2 O 2 – CaCl 2 – Ca 3 (PO 4 ) 2 (Ca 3 P 2 O 8 ) – N 2 O 4 – NO 2

18. Empirical Practice ● 56.79% Carbon, 6.56% Hydrogen, 28.37% Oxygen, 8.28% Nitrogen

19. Molecular Formula ● Molecular mass (given)/ empirical mass – Multiply empirical by the answer ● Example – SNH has a molecular mass of 188.35g/mol, what is the molecular formula? – CoC 4 O 4 has molecular mass of 341.94g/mol, what is the molecular formula?

20. Percent Composition ● Percent by mass – Element mass/compound mass x 100 ● NaCl ● Ca 3 (PO 4 ) 2

21. Combustion Analysis ● A 0.1156g sample composed of C, H, N is burned to produce 0.1638g CO 2 and 0.1676g H 2 O – Assume all C in the sample is converted to CO 2 – Assume all H is converted to H 2 O

22. Hydrates ● Ionic compound with water molecule attached in its crystal lattice – Example: ● CuSO 4 ·5H 2 O = copper(II)sulfate pentahydrate ● CuSO4 = anhydrous salt ● MgSO 4 · 7H 2 O =

23. Symbols ● Solid ● Liquid ● Gas ● Aqueous ● Reversible reaction ● Reactants → Products

24. Stoichiometry ● Use a balanced chemical equation and amount of at least one species to calculate amounts of other species ● Use mole ratios ● N 2 + H 2 → NH 3 ; how many moles of H 2 needed to react with 3mol N 2 ? ● How many molecules NH 3 produced from 13.5L N 2 ?

25. Stoichiometry ● CaO + NaNO 3 → Ca(NO 3 ) 2 + Na 2 O ● Mass of Na 2 O from 29.3g NaNO 3 ? ● Mass of Ca(NO 3 ) 2 from 2.93x10 25 molecules CaO? ● Mass of Ca(NO 3 ) 2 from 9.3mol NaNO 3 ?

26. Limiting Reactants ● Reactant that is “used up” in a chemical reaction ● Excess reactant – reactant that is left over ● Given amounts of more than one reactant ● 2 ways to solve: – Find the moles of each reactant, compare to moles in the reaction – Solve the problem twice (with each reactant), the reactant that produces the smallest amount is limiting

27. Why Use Excess Reactant? ● Can cause the reaction to go faster – Use the least expensive as the excess ● Can shift the reaction towards the products – Equilibrium reaction – reversible reaction

28. Yield ● Theoretical yield – calculated amount of product – Use stoichiometry ● Experimental yield – actual yield – Amount obtained in a lab – Would be given in a problem “obtained”, “recovered”, etc. ● Percent yield = experimental/theoretical x 100 ● Normally less than 100% – Incomplete reaction – Impure reagents – Experimental error – Side reactions