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Compounds & Moles

Compounds & Moles. Unit 5. Overview. Naming Ionic Covalent Acids Simple Organic The Mole Molar Mass Mole Conversions. Calculations Percent Composition Empirical Formula Molecular Formula. Why do we name compounds?. Think of some common compounds that you know of.

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Compounds & Moles

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  1. Compounds & Moles Unit 5

  2. Overview • Naming • Ionic • Covalent • Acids • Simple Organic • The Mole • Molar Mass • Mole Conversions • Calculations • Percent Composition • Empirical Formula • Molecular Formula

  3. Why do we name compounds? • Think of some common compounds that you know of • H2O = water NaCl = table salt • CaCO3 = limestone • Imagine if we had to memorize common names for the millions of known compounds that we had today • …IMPOSSIBLE! • Standard system was created to name compounds • IUPAC (International Union of Pure and Applied Chemistry)

  4. Chemical Formulas • Indicate the relative numbers of atoms of each kind in a chemical compound C8H18 Indicates 18 hydrogen atoms Indicates 8 carbon atoms

  5. Molecular vs. Structural Formulas • Molecular Formula • Lists elements in a compound and how many of each element you have • Example: C2H6O • Structural Formula • Shows how atoms are “connected” in the structure • Example CH3CH2OH or CH3OCH3

  6. Monatomic Ions • Ions formed from a single atom • Naming cations • Simply give the element’s name • Example • Ca+2 = calcium ion • Na+1 = sodium ion • Naming anions • Drop the ending of the element’s name and add “-ide” • Example • F-1 = fluoride ion • O-2 = oxide ion

  7. Binary Ionic Compounds • Ionic compound composed of 2 elements • Writing Names • Name the cation 1st • Name the anion 2nd • Example: • NaCl = sodium chloride • MgF2 = magnesium fluoride • Sr3N2 = strontium nitride

  8. Binary Ionic Compounds • Writing Formulas Example: aluminum oxide • Write the symbols for the ions side by side (cation first) Al+3 O-2 • Criss-cross the charges (use absolute value) Al2 O3 • Simplify (divide both numbers by largest common factor) Al2O3

  9. Binary Ionic Compounds • More examples (name to formula) • Calcium nitride = Ca3N2 • Potassium sulfide = K2S • Magnesium oxide = MgO

  10. Polyatomic Ions • Electrically charged group of two or more atoms • Oxyanion – polyatomic anion that contains oxygen • General naming rules • Most common oxyanion ends in “-ate” • Example • ClO3-1 = chlorate • NO3-1 = nitrate • SO4-2 = sulfate

  11. Polyatomic Ions • The number of oxygen atoms may be altered giving new endings and prefixes to oxyanions 1 more oxygen = per_______ate Common form = _______ate 1 less oxygen = _______ite 2 less oxygens = hypo_______ite • Example • ClO4-1 = perchlorate • ClO3-1 = chlorate • ClO2-1 = chlorite • ClO-1 = hypochlorite • Notice that the charge of the oxyanion does not change (only the number of oxygen atoms)

  12. Polyatomic Ions • Ionic compounds (contain “ions”) • Writing Name • If ion comes first, name the polyatomic ion then name the anion • If the ion comes second, name the cation then name the polyatomic ion (do not change ending) • Examples • NH4Cl = ammonium chloride • CaSO4 = calcium sulfate • Ba3(PO4)2 = barium phosphate

  13. Polyatomic Ions • Writing Formula • Follow same rules as binary ionic compound, but when charges are criss-crossed, use parenthesis to indicate number belongs to entire polyatomic ion • Example: calcium nitrate Ca+2 NO3-1 = Ca(NO3)2

  14. Stock System (Ionic Compounds) • For elements that form two or more cations with different charges (example Pb+2 and Pb+4) • Uses roman numeral to indicate ion’s charge • Transition metals, Sn, and Pb use this system • Writing Formulas • Roman numeral indicates charge of the cation (use that to criss cross) • Examples • Copper (II) bromide = CuBr2 • Iron (III) sulfide = Fe2S3 • Tin (IV) phosphate = Sn3(PO4)4

  15. Stock System (Ionic Compounds) • Writing Names • Use the anion (known charge) that the cation is bonded to and solve for the charge of the cation • Total positive charge (from cation) must equal total negative charge (from anion) • Example: VF6 • Fluorine has a charge of -1 • There are six fluorines bonded to the vanadium • 6 × -1 = -6 so the charge of vanadium is 6 • Name = vanadium (VI) fluoride

  16. Stock System (Ionic Compounds) • Example 2: Sn3N2 • The charge of nitrogen is -3 • There are 2 nitrogen atoms • 2 × -3 = -6 • There are 3 tin atoms that add up to a charge of +6 • +6 ÷ 3 = -2 so each tin atom has a charge of +2 • Name = tin (II) nitride • Exception: some transition metals only have one charge (nickel, silver, zinc, etc.) so the roman numeral is omitted

  17. Prefixes Used in naming covalent compounds Indicate how many of each atom you have

  18. Binary Covalent Compounds • Writing Names • Name the cation followed by the anion (-ide ending) • Use prefixes to indicate how many of each atom you have • Examples: • P4Br10 = tetraphosphorousdecabromide • Si2O5 = disiliconpentoxide • Note • If an o or a are doubled, drop the o or a of the prefix • Never use mono- on cation (only on anion)

  19. Binary Covalent Compounds • Writing Formulas • Prefix indicates how many of each atom you have • Do not criss-cross numbers • Examples: • Trinitrogenoctachloride = N3Cl8 • Arsenic tetrabromide = AsBr4

  20. Summary When writing names of formulas… YES NO YES NO

  21. Acids • Binary acid – contains two elements (one usually hydrogen and the other usually a halogen) • Oxyacid – acids that contain hydrogen, oxygen, and a third element (usually a nonmetal) • Usually hydrogen and a polyatomic ion

  22. Acids • Naming binary acids • Use form of hydro_____ic acid • Examples: • HF = hydrofluoric acid • HCl = hydrochloric acid

  23. Acids • As the number of oxygen atoms changes in oxyacids, so does the name (just like the oxyanions) 1 more oxygen = per_______ic acid Common form = _______ic acid 1 less oxygen = _______ous acid 2 less oxygens = hypo_______ous acid • Example • HClO4 = perchloric acid • HClO3 = chloric acid • HClO2 = chlorous acid • HClO = hypochlorous acid

  24. Carbon • Basis for all life. • Study of carbon compounds is called organic chemistry. • Can form single, double and triple bonds. • Long carbon chains can be produced. • Will bond with many other elements. • A HUGE number of compounds is possible (organic compounds)

  25. Naming Simple Organic Compounds • Organic compounds containing only carbon and hydrogen are called hydrocarbons • Alkane – all carbons form single bonds • Alkene – carbons form double bonds • Alkyne – carbons form triple bonds • Whether a compound is an alkane, alkene, or alkyne determines the suffix (ending) in the name of the hydrocarbon

  26. Number of carbons determines prefix used in name Naming Simple Organic Compounds • Prefix Carbons • Meth- 1 • Eth- 2 • Prop- 3 • But- 4 • Pent- 5 • Hex- 6 • Hept- 7 • Oct- 8 • Non- 9 • Dec- 10

  27. Naming Simple Organic Compounds • Examples • CH4 = methane • C2H6 = ethane propane propene propyne

  28. The Mole • The amount of a substance that contains as many particles as there are atoms in exactly 12 g of 12C • SI unit of amount of a substance • Abbreviated “mol” • Counting unit just like a “dozen” • 1 dozen donuts is the same amount as 1 dozen books • 1 mole of hydrogen atoms is the same amount as 1 mole of sodium atoms

  29. Avogadro’s Number • 6.022×1023 of anything is a mole • Named after Italian scientist Amadeo Avogadro • Experimentally determined number of atoms in 12 grams of 12C • How big is 602,200,000,000,000,000,000,000? • One mole of donut holes would cover the Earth 5 miles deep in the donut holes • One mole of pennies stacked on top of each other would reach from the Earth to the moon 7 times • If you started counting when you were born and never stopped until the day you died, you would never come close to reaching 6.022×1023

  30. Avogadro’s Number • 1 Liter of water contains 55.5 moles of H2O • A 5 lb bag of sugar contains 6.6 moles of sugar How can that be?! • Atoms and molecules are so tiny that when we use units of moles (6.022×1023) it puts the particles into measurable quantities

  31. Molar Mass • 1 mole of hydrogen atoms = 1 mole of sodium atoms BUT… • 1 mole of hydrogen atoms DOES NOT have the same mass as 1 mole of sodium atoms • Individual atoms have different masses • They are the same amount but not the same mass

  32. Molar Mass • The periodic table tells us the mass of 1 mole of any atom • It’s the same as the average atomic mass/relative atomic mass (decimal number on the table) • Molar Mass – mass of 1 mole of an atom or compound • Units are “grams/mole” or “g/mol”

  33. Molar Mass • To find the molar mass of a compound, add the molar masses of all atoms in a compound • Also called formula mass or molecular mass (compounds only) Example: CO2 (1 atom of C and 2 atoms of O) 1 atom C x 12.011 = 12.011 2 atoms O x 15.9994 = 31.9988 Molar mass = 44.010 g/mol

  34. Mole Relationships • To go between units of grams, moles and atoms (or molecules) use conversions! • 6.022×1023 is how many atoms or molecules are in 1 mole of any substance • The molar mass is how many grams are in one mole of any substance 6.02 x 1023 Molar Mass

  35. Mole Conversions • How many grams are in 5.0 moles of calcium? 5.0 mole × = 200.39 g • How many atoms are in 2.1 moles of xenon? 2.1 moles × = 1.26×1024 atoms 40.078 g 1 mole 6.022×1023 atoms 1 mole

  36. Mole Conversions • There is no way to go straight from grams to atoms or molecules in one step • Must use moles as the intermediate step • How many atoms are in 9.8 g of Pb? 9.8g × × = 2.8×1022 atoms 1 mol 207.2 g 6.022×1023 atoms 1 mole

  37. Mole Conversions • When a conversion includes a compound, it will use the word molecules when a conversion includes an element, it will use the word atoms • There are still as many molecules in a mole as there are atoms • How many grams are in 3.4×1022 molecules of H2O? • First solve for molar mass of H2O (H2O molar mass = 18.02g/mol) 3.4×1022 molecules × × = 1.0 g 18.02 g 1 mole 1 mole 6.022×1023 molecules

  38. Percent Composition • Percentage by mass of each element in a compound Example: What is the percent composition of BaSO4? Ba = 1 × 137.3 = 137.3 (137.3/233.4) ×100= 58.8% Ba S = 1 × 32.1 = 32.1 (32.1/233.4) ×100= 13.8% S O = 4 × 16.0 = 64.0 (64.0/233.4) ×100= 27.4% O 233.4 Multiply by 100 Molar Mass part ÷ total Total molar mass

  39. Empirical Formula • Smallest whole-number ratio formula of a compound • Simplest formula • What is the empirical formula of a compound that is 27.0% sodium, 16.5% nitrogen, and 56.5% oxygen by mass? • Assume that you have a 100 gram sample Na 27.0/22.99 = 1.17 /1.17 = 1 N 16.5/14.01 = 1.18 /1.17 = 1 O 56.5/16.00 = 3.53 /1.17 = 3 Divide by smallest number Molar Mass Empirical Formula = NaNO3

  40. Empirical Formula • When numbers are too far to round, you may need to multiply all values by the same factor to make all numbers whole What is the empirical formula of a compound that contains 40.6g of calcium and 9.5g of nitrogen? Ca 40.6/40.1 = 1.01 /0.69 = 1.5 × 2 = 3 N 9.5/14.01 = 0.69 /0.69 = 1 × 2 = 2 too far to round double both numbers to get whole numbers Empirical Formula = Ca3N2

  41. Molecular Formula • Indicates actual number of atoms of each element in a compound • Multiple of empirical formula • An empirical formula can be the molecular formula, but the molecular formula is not always the empirical formula Molecular Formula Empirical Formula C3H12 CH4

  42. Molecular Formula • If the molecular mass is known, you can solve for the molecular formula • The molar mass of a compound with empirical formula of CH2O is 180.12 g/mol. What is the molecular formula of this compound? Molar mass CH2O = 30.02g/mol = 6 Molecular Formula = CH2O × 6 = C6H12O6 180.12 30.02

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