Chapter 3 PowerPoint - Home - Faculty · • The attraction between cation and anion. • Atoms of...
Transcript of Chapter 3 PowerPoint - Home - Faculty · • The attraction between cation and anion. • Atoms of...
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Chemistry 1AChapter 3
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Covalent Bond Formation
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Covalent Bond• A link between atoms due to the sharing of
two electrons. This bond forms between atoms of two nonmetallic elements.– If the electrons are shared equally, there is a
even distribution of the negative charge for the electrons in the bond, so there is no partial charges on the atoms. The bond is called a nonpolar covalent bond.
– If one atom in the bond attracts electrons more than the other atom, the electron negative charge shifts to that atom giving it a partial negative charge. The other atom loses negative charge giving it a partial positive charge. The bond is called a polar covalent bond.
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Polar Covalent Bond
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Ionic Bond• The attraction between cation and anion. • Atoms of nonmetallic elements often attract
electrons so much more strongly than atoms of metallic elements that one or more electrons are transferred from the metallic atom (forming a positively charged particle or cation), to the nonmetallic atom (forming a negatively charged particle or anion).
• For example, an uncharged chlorine atom can pull one electron from an uncharged sodium atom, yielding Cl− and Na+.
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Ionic Bond Formation
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Sodium Chloride, NaCl, Structure
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Electronegativities
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Electronegativity, a measure of the electron attracting ability of atoms in chemical bonds is used to predict…
• whether a chemical bond is nonpolar covalent, polar covalent, or ionic.
• which atom in a polar covalent bond is partial negative and which is partial positive.
• which atom in an ionic bond forms the cation and which forms the anion.
• which of two covalent bonds are more polar.
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Bond Types
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Which atom in a polar covalent bond is partially negative and which is partially positive?
higher electronegativity↓
partial negative charge
lower electronegativity↓
partial positive charge
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Which of two bonds is more polar?
The greater the ΔEN is, the more polar the bond.
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Types of Compounds
• All nonmetallic atoms usually leads to all covalent bonds, which from molecules. These compounds are called molecular compounds.
• Metal-nonmetal combinations usually lead to ionic bonds and ioniccompounds.
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Classification of Compounds
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Summary• Nonmetal-nonmetal combinations
(e.g. HCl)– Covalent bonds– Molecules– Molecular Compound
• Metal-nonmetal combinations (e.g. NaCl)– Probably ionic bonds– Alternating cations and anions in
crystal structure– Ionic compound
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Valence Electrons
• The valence electrons for each atom are the most important electrons in the formation of chemical bonds.
• The number of valence electrons for the atoms of each element is equal to the element’s A-group number on the periodic table.
• Covalent bonds often form to pair unpaired electrons and give the atoms of the elements other than hydrogen and boron eight valence electrons (an octet of valence electrons).
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Valence Electrons and A-Group Numbers
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Electron-Dot Symbols and Lewis Structures
• Electron-dot symbols show valence electrons.
• Nonbonding pairs of valence electrons are called lone pairs.
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Lewis Structures
• Lewis structures represent molecules using element symbols, lines for bonds, and dots for lone pairs.
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Most Common Bonding Patterns for Nonmetals
Element Number of bonds
Number of lone pairs
H 1 0
C 4 0
N, P 3 1
O, S, Se 2 2
F, Cl, Br, I 1 3
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Drawing Lewis Structures
• A later chapter describes procedure that allows you to draw Lewis structures for many different molecules.
• Many Lewis structures can be drawn by attempting to give each atom in a molecule its most common bonding pattern.
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Lewis Structure for Methane, CH4
• Carbon atoms usually have 4 bonds and no lone pairs.
• Hydrogen atoms have 1 bond and no lone pairs.
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Tetrahedral Geometry
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Methane, CH4
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Lewis Structure for Ammonia, NH3
• Nitrogen atoms usually have 3 bonds and 1 lone pair.
• Hydrogen atoms have 1 bond and no lone pairs.
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Ammonia, NH3
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Lewis Structure for Water, H2O
• Oxygen atoms usually have 2 bonds and 2 lone pairs.
• Hydrogen atoms have 1 bond and no lone pairs.
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Water, H2O
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Water Attractions
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Liquid Water
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Study Sheets for Converting Between Names and Formulas
• For each type of compound,– How recognize from formula– How recognize from name– How write formula– How write name
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Binary Covalent
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Common Names
–H2O, water–NH3, ammonia–CH4, methane –C2H6, ethane–C3H8, propane
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Naming Binary Covalent Compounds
• If the subscript for the first element is greater than one, indicate the subscript with a prefix.– We do not write mono- on the first name.– Leave the "a" off the end of the prefixes that
end in "a" and the “o” off of mono- if they are placed in front of an element that begins with a vowel (oxygen or iodine).
• Follow the prefix with the name of the first element in the formula.
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Naming Binary Covalent Compounds
• Write a prefix to indicate the subscript for the second element.
• Write the root of the name of the second symbol in the formula.
• Add -ide to the end of the name.
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Prefixes
mon(o)ditritetr(a)pent(a)
hex(a)hept(a)oct(a)non(a)dec(a)
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Roots of Nonmetals
H hydr-C carb-N nitr-P phosph-O ox-S sulf-Se selen-
F fluor-Cl chlor-Br brom-I iod-
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Binary Covalent Names without Prefixes
Formula Complete name Abbreviated Name
HF Hydrogen monofluoride
Hydrogen fluoride
HCl Hydrogen monochloride
Hydrogen chloride
HBr Hydrogen monobromide
Hydrogen bromide
HI Hydrogen moniodide Hydrogen iodide
H2S Dihydrogen monosulfide
Hydrogen sulfide
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The Making of an Anion
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The Making of a Cation
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Monatomic Ions
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Sodium Chloride, NaCl, Structure
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Monatomic Ion Names
• Monatomic Cations– (name of metal)
• Groups 1, 2, and 3 metals• Al3+, Zn2+, Cd2+, Ag+
– (name of metal)(Roman numeral)• All metallic cations not mentioned above
• Monatomic Anions– (root of nonmetal name)ide
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Naming Cations with Two Possible Charges
Ion Systematic Name
Nonsystematic Name
Fe2+ iron(II) ferrous
Fe3+ iron(III) ferric
Cu+ copper(I) cuprous
Cu2+ copper(II) cupric
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Hydride H−
Nitride N3−
Phosphide P3−
Oxide O2−
Sulfide S2−
selenide Se2−
fluoride F−
chloride Cl−
bromide Br−
iodide I−
Monatomic Anions
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CsCl and NH4Cl structure
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Polyatomic Ions
Ion Name Ion Name
NH4+ ammonium NO3
− nitrate
OH− hydroxide SO42− sulfate
CO32− carbonate C2H3O2
− acetate
PO43− phosphate
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(Root)ate Polyatomic Ions
Ion Name Ion Name
SO42− sulfate NO3
− nitrate
PO43− phosphate CO3
2− carbonate
ClO3− chlorate CrO4
2− chromate
BrO3− bromate IO3
− iodate
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Convention for Naming Oxyanions
Relationship General Name
Example Name
Example Formula
One more O per(root)ate perchlorate ClO4−
(root)ate chlorate ClO3−
One less O (root)ite chlorite ClO2−
Two less O hypo(root)ite hypochlorite ClO−
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Polyatomic Ions with Hydrogen
• HCO3− hydrogen carbonate
• HSO4− hydrogen sulfate
• HSO3− hydrogen sulfite
• HS− hydrogen sulfide• HPO4
2− hydrogen phosphate• H2PO4
− dihydrogen phosphate
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Systematic and Nonsystematic Names
Formula Systematic Name Nonsystematic Name
HCO3− hydrogen carbonate bicarbonate
HSO4− hydrogen sulfate bisulfate
HSO3− hydrogen sulfite bisulfite
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Study Sheets for Converting Between Names and Formulas
• For each type of compound,– How recognize from formula– How recognize from name– How write formula– How write name
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Names and Formulas of Ionic Compounds
• Name– (name of cation) (name of anion)
• Formula Steps– Determine the Formula, including
charge, for the cation and anion. – Determine the ratio of the ions that
yields zero overall charge.
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Cation Names
Metals with one possible charge (Al,
Zn, Cd, and Groups 1, 2, 3)
name of metal
Metals with more than one possible charge
(the rest)
name(Roman numeral)
polyatomic cations (e.g. ammonium)
name of polyatomic ion
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Anion Names
monatomic anion (root of nonmetal name)ide
polyatomic anion name of polyatomic ion
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Binary Acid Names and Formulas
• General – HX(aq) hydro(root)ic acid– HF(aq) hydrofluoric acid– HCl(aq) hydrochloric acid– HBr(aq) hydrobromic acid– HI(aq) hydroiodic acid– H2S(aq) hydrosulfuric acid
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Oxyacid Names and Formulas
• When enough H+ ions are added to a (root)ate polyatomic ion to neutralize its charge, the (root)ic acid.– nitrate, NO3
−, forms nitric acid, HNO3
– sulfate, SO42−, forms sulfuric acid, H2SO4
– phosphate, PO43−, forms phosphoric acid, H3PO4
– carbonate, CO32−, forms carbonic acid, H2CO3
– chlorate, ClO3−, forms chloric acid, HClO3
– bromate, BrO3−, forms bromic acid, HBrO3
– iodate, IO3−, forms iodic acid, HIO3
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Oxyacid Names and Formulas (cont)
Relationship General Name
Example Name
Example Formula
One more O per(root)ic acid perchloric acid HClO4
(root)ic acid chloric acid HClO3
One less O (root)ous acid chlorous acid HClO2
Two less O hypo(root)ous acid
hypochlorous acid HClO
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Common Alcohols and Sugars
• CH3OH methanol (methyl alcohol)
• C2H5OH ethanol (ethyl alcohol)
• C3H7OH 2-propanol (isopropyl alcohol)
• C6H12O6 glucose (or fructose or galactose)
• C12H22O11 sucrose(or maltose or lactose)
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Condensation (Gas to Liquid)
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Types of Particles
Type of substance
Particles to visualize
Type of substance
Particles to visualize
metalscations in a
sea of electrons
ionic cations and anions
noble gases atoms molecular molecules
carbon (diamond) atoms
other nonmetallic elements
molecules
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Types of Attractions –Metallic Elements
• Metallic bonds = the attraction between the positive metal cations and the negative electrons that surround them.
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Types of Attractions –Ionic Compounds
• Cations and Anions held together by ionic bonds
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Types of Attractions –Molecular Compounds
• Nonpolar molecular compounds (e.g. hydrocarbons, CaHb) – London forces
• Polar molecular compounds without H-F, O-H, or N-H (e.g. HCl) – dipole-dipole attractions
• Polar molecular compounds with H-F, O-H, or N-H (e.g. HF, H2O, NH3, alcohols, and sugars) – hydrogen bonds
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Dipole-Dipole Attractions
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Dipole-Dipole Attractions in a Liquid
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Hydrogen Bonds
• Hydrogen bond = the intermolecular attraction between the partially negative O, N, or F of one molecule and a partially positive H connected to O, N, or F of another molecule.
• Relatively strong
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Hydrogen Bonds in HFIn HF, the hydrogen bond is between the partial positive H of one HF molecule and the partial negative F of another HF molecule.
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Hydrogen Bonds in WaterIn H2O, the hydrogen bond is between the partial positive H of one H2O molecule and the partial negative O of another H2O molecule.
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Hydrogen Bonds in MethanolIn CH3OH, the hydrogen bond is between the partial positive H of one CH3OH molecule and the partial negative O of another CH3OH molecule.
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Hydrogen Bonds in AmmoniaIn NH3, the hydrogen bond is between the partial positive H of one NH3 molecule and the partial negative N of another NH3molecule.
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London Forces
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London Forces in Polar Molecules
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Why Larger Molecules Have Stronger London Forces
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Types of Attractions –Carbon
• Diamond - Carbons atoms held together by covalent bonds, forming huge 3-dimensional molecules.
• Graphite - Carbons atoms held together by covalent bonds, forming huge 2-dimensional molecules held together by London forces.
• Fullerenes - Carbons atoms held together by covalent bonds, forming 3-dimensional molecules held together by London forces.
• See the following website.
http://preparatorychemistry.com/Bishop_Jmol_Carbon.htm
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Predicting Types of Attractions
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Types of Particles and Attractions - ElementsType of element
Particles to visualize
Examples Type of attraction
metals cations in a sea of electrons gold, Au metallic
bonds
noble gases atoms xenon, Xe London forces
carbon (diamond) atoms C(dia) covalent
bonds
other nonmetallic elements
molecules H2, N2, O2, F2, Cl2, Br2, I2, S8, Se8, P4
London forces
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Types of Particles and Attractions - Compounds
Type of compound Particles to visualize
Examples Type of attraction
ionic cations and anions NaCl ionic
bonds
nonpolar molecular molecules hydrocarbons London forces
polar molecular w/out H-F, O-H, or N-H molecules HCl dipole-
dipole
polar molecular with H-F, O-H, or N-H molecules HF, H2O, NH3
alcohols, hydrogen
bonds
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Making Phosphoric Acid
• Furnace Process for making H3PO4 to be used to make fertilizers, detergents, and pharmaceuticals. – React phosphate rock with sand and coke at 2000
°C.2Ca3(PO4)2 + 6SiO2 + 10C
→ 4P + 10CO + 6CaSiO3– React phosphorus with oxygen to get
tetraphosphorus decoxide.4P + 5O2 → P4O10
– React tetraphosphorus decoxide with water to make phosphoric acid.
P4O10 + 6H2O → 4H3PO4
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Sample Calculations (1)
• What is the maximum mass of P4O10 that can be formed from 1.09 × 104 kg P?
• The formula for P4O10 provides us with a conversion factor that converts from units of P to units of P4O10.
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Sample Calculations (2)• What is the minimum mass of water that
must be added to 2.50 × 104 kg P4O10 to form phosphoric acid in the following reaction?
P4O10 + 6H2O → 4H3PO4• The coefficients in the balanced
equation provide us with a conversion factor that converts from units of P4O10to units of H2O.
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Goal: To develop conversion factors that will convert between a measurable property (mass) and number of particles
Measurable Property 1↓
Number of Particles 1↓
Number of Particles 2↓
Measurable Property 2
Mass 1↓
Number of Particles 1↓
Number of Particles 2↓
Mass 2
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Counting by Weighing for Nails
• Step 1: Choose an easily measurable property.– Mass for nails
• Step 2: Choose a convenient unit for measurement.– Pounds for nails
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Counting by Weighing for Nails (cont)
• Step 3: If the measurable property is mass, determine the mass of the individual objects being measured.– Weigh 100 nails: 82 are 3.80 g, 14 are 3.70
g, and 4 are 3.60 g
• Step 4: If the objects do not all have the same mass, determine the weighted average mass of the objects.
0.82(3.80 g) + 0.14(3.70 g) + 0.04(3.60 g) = 3.78 g
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Counting by Weighing for Nails (cont)
• Step 5: Use the conversion factor from the weighted average to make conversions between mass and number of objects.
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Counting by Weighing for Nails (cont)
• Step 6: Describe the number of objects in terms of a collective unit such as a dozen, a gross, or a ream.
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Counting by Weighing for Carbon Atoms
• Step 1: Choose an easily measurable property.– Mass for carbon atoms
• Step 2: Choose a convenient unit for measurement.– Atomic mass units (u) for carbon atoms– Atomic mass unit (u) = 1/12 the mass of a
carbon-12 atom (with 6 p, 6 n, and 6 e−)
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Counting by Weighing for Carbon Atoms (cont)
• Step 3: If the measurable property is mass, determine the mass of the individual objects being measured.– For carbon: 98.90% are 12 u and 1.10% are
13.003355
• Step 4: If the objects do not all have the same mass, determine the weighted average mass of the objects.
0.9890(12 u) + 0.0110(13.003355 u) = 12.011 u
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Mass Spectrometer
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Mass Spectrum
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Counting by Weighing for Carbon Atoms (cont)
• Step 5: Describe the number of objects in terms of a collective unit such as a dozen, a gross, or a ream, and use this and the weighted average to create a conversion factor to make conversions between mass and number of objects.
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Mole• A mole (mol) is an amount of substance
that contains the same number of particles as there are atoms in 12 g of carbon-12.
• To four significant figures, there are 6.022 × 1023 atoms in 12 g of carbon-12.
• Thus a mole of natural carbon is the amount of carbon that contains 6.022 × 1023 carbon atoms.
• The number 6.022 × 1023 is often called Avogadro’s number.
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Avogadro’s Number
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Molar Mass Development
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Molar Mass For Elements
• Atomic Mass from the Periodic Table
⎛ ⎞⎜ ⎟⎝ ⎠
(atomic mass) g element1 mol element
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Molar Mass Calculation for Carbon
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Goal: To develop conversion factors that will convert between a measurable property (mass) and number of particles
Measurable Property 1↓
Number of Particles 1↓
Number of Particles 2↓
Measurable Property 2
Mass 1↓
Moles 1↓
Moles 2↓
Mass 2
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Molecular Mass
• Whole = sum of parts• mass of a molecule = sum of the
masses of the atoms in the molecule • molecular mass = the sum of the
atomic masses of the atoms in the molecule
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Molar Mass For Molecular Compounds
• Molecular Mass = Sum of the atomic masses of atoms in one molecule
⎛ ⎞⎜ ⎟⎝ ⎠
(molecular mass) g molecular compound1 mol molecular compound
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Formula Units
• A formula unit of a substance is the group represented by the substance’s chemical formula, that is, a group containing the kinds and numbers of atoms or ions listed in the chemical formula.
• Formula unit is a general term that can be used in reference to elements, molecular compounds, or ionic compounds.
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Formula Unit Examples
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Formula Mass for Ionic Compounds
• Whole = sum of parts• Mass of a formula unit = sum of the
masses of the atoms in the formula unit • Formula mass = the sum of the atomic
masses of the atoms in the formula
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Molar Mass For Ionic Compounds
• Formula Mass = Sum of the atomic masses of the atoms in a formula unit
⎛ ⎞⎜ ⎟⎝ ⎠
(formula mass) g ionic compound1 mol ionic compound
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Molar Mass Development
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General Conversions
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Units of One Substance to Units of Another
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Study Sheets
• Write a description of the “tip-off” that helps you to recognize the type of problem the calculation represents.
• Write a description of the general procedure involved in the particular type of problem.
• Write an example of the type of calculation.
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Sample Study Sheet: Converting Between Mass of Element and Mass of Compound Containing the Element
• Tip-off: When you analyze the type of unit you have and the type of unit you want, you recognize that you are converting between a unit associated with an element and a unit associated with a compound containing that element.
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Sample Study Sheet (2)
• General Steps– Convert the given unit to moles of the
first substance.– Convert moles of the first substance
to moles of the second substance using the molar ratio derived from the formula for the compound.
– Convert moles of the second substance to the desired units of the second substance.
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Units of Element
to Units of Compound
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Beginning and End