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COMPOUNDS & MOLES Unit 5. Overview  Naming  Ionic  Covalent  Acids  Simple Organic  The Mole  Molar Mass  Mole Conversions  Calculations  Percent.

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Presentation on theme: "COMPOUNDS & MOLES Unit 5. Overview  Naming  Ionic  Covalent  Acids  Simple Organic  The Mole  Molar Mass  Mole Conversions  Calculations  Percent."— Presentation transcript:

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  H 2 O = waterNaCl = table salt  CaCO 3 = 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 C 8 H 18 Indicates 8 carbon atoms Indicates 18 hydrogen atoms

5 Molecular vs. Structural Formulas  Molecular Formula  Lists elements in a compound and how many of each element you have  Example: C 2 H 6 O  Structural Formula  Shows how atoms are “connected” in the structure  Example CH 3 CH 2 OH or CH 3 OCH 3

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 2 nd  Example: NaCl = sodium chloride MgF 2 = magnesium fluoride Sr 3 N 2 = 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) Al 2 O 3  Simplify (divide both numbers by largest common factor) Al 2 O 3

9 Binary Ionic Compounds  More examples (name to formula)  Calcium nitride = Ca 3 N 2  Potassium sulfide = K 2 S  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 ClO 3 -1 = chlorate NO 3 -1 = nitrate SO 4 -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  ClO 4 -1 = perchlorate  ClO 3 -1 = chlorate  ClO 2 -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 NH 4 Cl = ammonium chloride CaSO 4 = calcium sulfate Ba 3 (PO 4 ) 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 NO 3 -1 = Ca(NO 3 ) 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 = CuBr 2 Iron (III) sulfide = Fe 2 S 3 Tin (IV) phosphate = Sn 3 (PO 4 ) 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: VF 6  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: Sn 3 N 2  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 NumberPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

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: P 4 Br 10 = tetraphosphorous decabromide Si 2 O 5 = disilicon pentoxide  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: Trinitrogen octachloride = N 3 Cl 8 Arsenic tetrabromide = AsBr 4

20 Summary Is it ionic? Is the cation a transition metal, Sn, or Pb? Use Roman numerals Name cation then anion (write it like it is) Use prefixes (covalent) NOYES NO When writing names of formulas…

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  HClO 4 = perchloric acid  HClO 3 = chloric acid  HClO 2 = 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. HUGE 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 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 Number of carbons determines prefix used in name

27 Naming Simple Organic Compounds  Examples  CH 4 = methane  C 2 H 6 = ethane propanepropenepropyne

28 The Mole  The amount of a substance that contains as many particles as there are atoms in exactly 12 g of 12 C  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  6.022×10 23 of anything is a mole  Named after Italian scientist Amadeo Avogadro  Experimentally determined number of atoms in 12 grams of 12 C  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 6.022×10 23

30 Avogadro’s Number  1 Liter of water contains 55.5 moles of H 2 O  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 ( 6.022×10 23 ) 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: CO 2 (1 atom of C and 2 atoms of O) 1 atom C x = atoms O x = Molar mass = g/mol Molar mass = g/mol

34 Atoms Molecules Mole Grams Mole Relationships 6.02 x Molar Mass  To go between units of grams, moles and atoms (or molecules) use conversions!  6.022×10 23 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

35 Mole Conversions  How many grams are in 5.0 moles of calcium? 5.0 mole × = g  How many atoms are in 2.1 moles of xenon? 2.1 moles × = 1.26×10 24 atoms g 1 mole 6.022×10 23 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 × × = 2.8×10 22 atoms 1 mol g 6.022×10 23 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 3.4×10 22 molecules of H 2 O?  First solve for molar mass of H 2 O (H 2 O molar mass = 18.02g/mol) 3.4×10 22 molecules × × = 1.0 g 1 mole 6.022×10 23 molecules g 1 mole

38 Percent Composition  Percentage by mass of each element in a compound Example: What is the percent composition of BaSO 4 ? Ba = 1 × = 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 Molar Mass part ÷ total Multiply by 100 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 Empirical Formula = NaNO 3 Molar Mass Divide by smallest number

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 Empirical Formula = Ca 3 N 2 too far to round double both numbers to get whole numbers

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 CH 4 Empirical Formula C 3 H 12 Molecular Formula

42  If the molecular mass is known, you can solve for the molecular formula  The molar mass of a compound with empirical formula of CH 2 O is g/mol. What is the molecular formula of this compound? Molar mass CH 2 O = 30.02g/mol = 6 Molecular Formula = CH 2 O × 6 = C 6 H 12 O


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