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Compounds & Moles Unit 5.

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Presentation on theme: "Compounds & Moles Unit 5."— Presentation transcript:

1 Compounds & Moles Unit 5

2 Overview Naming The Mole Calculations 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 Indicates 18 hydrogen atoms
Chemical Formulas Indicate the relative numbers of atoms of each kind in a chemical compound C8H18 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: 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” 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 NO = 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 Number Prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca- 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 = tetraphosphorous decabromide Si2O5 = 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 = N3Cl8 Arsenic tetrabromide = AsBr4

20 Summary When writing names of formulas… YES NO YES NO 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) 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 Naming Simple Organic Compounds
Number of carbons determines prefix used in name Prefix Carbons Meth- 1 Eth- 2 Prop- 3 But- 4 Pent- 5 Hex- 6 Hept- 7 Oct- 8 Non- 9 Dec

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 How can that be?!
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 = 2 atoms O x = Molar mass = g/mol

34 Mole Relationships Mole Grams Atoms Molecules
6.02 x 1023 Molar Mass 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

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 × = 1.26×1024 atoms 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 1 mole 6.022×1023 molecules 18.02 g 1 mole

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

39 Divide by smallest number
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 N / = /1.17 = 1 O / = /1.17 = 3 Divide by smallest number Molar Mass Empirical Formula = NaNO3

40 Empirical Formula = Ca3N2
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 = /0.69 = × 2 = 3 N / = /0.69 = × 2 = 2 too far to round double both numbers to get whole numbers Empirical Formula = Ca3N2

41 CH4 C3H12 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 Empirical Formula Molecular Formula CH4 C3H12

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 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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