Chemical Bonds, Nomenclature, Lewis Structure and Molecular Shapes

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Chapter 5

• Chemical Bonds, Nomenclature, Lewis Structure and
  Molecular Shapes

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Homework Quizzes Chapter 5

• Text Homework (not turned in) pages 147 151.
  Problems 1, 6 8, 17, 27, 35, 39, 41 66, 68,
  69, 72, 86, 88(bc), 107, 109, 110, 112, 113.
• Quiz Do the graded quiz in Blackboard.

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I. Chemical Bonds A. Introduction (summary of
chapter)

• Atoms can combine to produce new larger units
  called molecules or compounds.
• Each molecule has a unique name (two rules to
  learn).
• Molecules held together by chemical bonds (two
  types).
• Bonds result from either transfer of valence
  electrons (Ionic Bonds) or from sharing of
  valence electrons Covalent Bonds.
• Valence electrons rearrange to mimic closest
  Group VIIIA (18) structure.
• Molecules resulting from covalent bonding will
  have predictable shapes.

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I. Chemical Bonds B. Ionic Bonds

• - Metals (except H) loose electrons, form
  cations Nonmetals gain electrons to form anions.
  Both strive for e- configuration of nearest
  inert gas.
• - The resulting opposite ions attract in a ratio
  which produces a neutral unit. Reduce formula to
  simplest ratio.
• Ionic Bond Definition bond formed by
  electrostatic attraction between anions (-) and
  cations ().
• Write formula with element first do not show
  charges Final compound is neutral.
• Generality any compound formed from metallic
  and nonmetallic elements is ionic.

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I. Chemical Bonds B. Ionic Bonds

• Know Metals combine with nonmetals form ionic
  bonds by losing or gaining electrons to mimic
  closest Inert Gas (VIIIA).
• IA - Na, K, Li, etc become 1 ions
  Na
• IIA - Ca, Mg, etc become 2 ions Ca2
• IIIA - Al, Ga become 3 ions Al3
• VA - N, P become -3 ions N-3
• VIA - O, S become -2 ions O-2
• VIIA - F, Cl, Br, I become -1 ions F-1
• Opposite ions attract in a ratio so that the
  product is neutral.

Inert Gas e- Configurations
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I. Chemical Bonds B. Ionic Bonds Example
Do not show charges in final formula. NaCl NOT
NaCl-
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I. Chemical Bonds B. Ionic Bonds Example
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I. Chemical Bonds B. Ionic Bonds - Examples

• Na Na O-2 Na2O
• Ca2 F- F- CaF2
• Mg2 S-2 MgS
• Al3 Al3 O-2 O-2 O-2 Al2O3
• Give the formulas for the following Na Br
• Ca O Ba I Li O Al F Mg
  N
• Many transition metals form ionic bonds can
  have several charges such as Fe2 Iron (II)
  Fe3 Iron (III) Cu2 Copper (II) Cu1
  Copper (I)

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I. Chemical Bonds C. Electron Dot (Lewis)
Structures

• - A Lewis electron dot structure is a symbol in
  which the valence electrons are shown as dots.
• - Examples Na. Mg Na
  Ca2
• H1- (Called Hydride) C Si
• - How many valence electrons (dots) would
• N3- O2- F- or Ne
  have? What about Mg2?
• 8 8 8 8
  0

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II. Covalent Bonds A. Introduction EN
electronegativity

• - Definition of a covalent bond A bond formed
  by the sharing of two electrons.
• When two atoms of similar EN combine, neither has
  the pull to take electrons away a sharing of
  electrons results.
• This occurs when NONMETALS, including H, combine
  with NONMETALS.
• Example H. H. ---) HH H2
• - The atoms share valence electrons to get stable
  group VIIIA e- configurations.

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II. Covalent Bonds A. Introduction

• Covalent bond sharing of 2 electrons.
• 2 shared electrons with (Single Bond).
• 4 shared electrons with (Double
  Bond).
• 6 shared electrons with (Triple
  Bond).
• We frequently show the structure as a Lewis
  Structure - covalent bonds with lines and
  nonbonding valence electrons as dots.
• - Note Group IVA usually forms 4 bonds VA
  three bonds VIA two bonds and VIIA (along with
  H) one bond.

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II. Covalent Bonds B. Examples

• H. F ---) H F
• H. O .H ---) H O H
• N N ---) N N
• Cl. .O. .Cl ---) Cl O
  Cl
• O C O -----) O C O

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II. Covalent Bonds C. Lewis Structures
1. Rules for Drawing Lewis Structures

• 1. Calculate the total of valence electrons
  take into account charge if the sample is an ion.
• 2. Place atom that forms most bonds at center
  (Closest to Group IVA Lowest if in same group).
  If there is a charge, then add or subtract the
  appropriate number of electrons on the central
  atom.
• 3. Arrange other atoms around central atom
  allow sharing so that each atom has stable
  electron configuration. Show bonding pairs as
  dashes nonbonding valence e- as dots.
• 4. Double check a) each atom has a stable
  electron configuration b) have the same total
  number of valence electrons as in step 1.

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II. Covalent Bonds C. Lewis Structures
2. Examples

• H I H2O NH4
• H2O2 CH4 SO2
• AlCl4- NO2- CN-

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Bonding SummaryTwo General Bonding Types

• Ionic Compound containing metallic element.
  Atoms lose/gain e to look like nearest inert gas.
  Add together ions such that neutralize charge.
• Ia IIa IIIa Va
  VIa VIIa
• 1 2 3 -3
  -2 -1
• 2. Covalent Compound containing nonmetals.
• Atoms obtain inert gas configuration by
• sharing valence electrons.
  

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II. Covalent Bonds Organic Compounds

• Can write organic structures several ways.
• Example Butane (Note the five ways of
  presenting)
• Note Carbon always has four bonds.
• C4H10 CH3CH2CH2CH3
  CH3-CH2-CH2-CH3
• H H H H
• H C C C C H
• H H H H

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II. Covalent Bonds Organic Compounds

• Cyclic Organics Example of Cyclopropane
• Aromatics Contain Benzene, C6H6

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II. Covalent Bonds Organic Compounds
C7H6O3 MW 138g
C9H8O4 MW 180 g
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II. Covalent Bonds Organic Compounds Aspirin
Lab

• 1) Equation Conversion Factors
• 1 Salicylic Acid 1 Acetic Anhydride -----)
  1 Aspirin 1 Acetic Acid
• 1 molecules or moles 1 mole formula
  weight in grams 6.0x1023 molecules
• 2) Lab Calculations (questions 2 3)
• 2.0 g SA x 1 mole SA 0.014 mole SA
• 138 g SA
• 0.014 mole SA x 1 mole Aspirin 0.014 mole
  Aspirin
• 1 mole SA

From the coefficients in the balanced chemical
equation above.
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III. Shapes

• Molecular Shapes play a major role in
• 1) Physical Properties
• 2) Chemical Properties
• 3) Biochemical Properties
• To Obtain the shape of a molecule one draws the
  Lewis Structure, counts the number of things
  around the central atom, and uses simple geometry
  to predict the shape.

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III. Shapes C. Simplified Examples
Bond angle 180o
Bond angle 120o
Bond angle 109o
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IV. Nomenclature A. Introductions

• There are common systematic names for
  chemicals. A chemical may have scores of common
  names.
• A systematic name must allow one to both obtain
  the formula and derive the name from the formula.
  
• There are two general rules for naming inorganic
  compounds.
• Ionic compounds use Rule 1. Molecular or
  Covalent compounds use Rule 2.

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IV. Nomenclature B. Ionic Compounds

• Rule 1 for ionic compounds Name the element,
  then the element and change the ending to
  ide.
• Examples
• NaCl Sodium Chloride
• Na2O Sodium Oxide

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IV. Nomenclature B. Ionic Compounds Rule 1
ide names

• Negative atoms have an ide ending.
• Atom Anion Name
• Chlorine Cl1- Chloride
• Oxygen O2- Oxide
• Fluorine F1- Fluoride
• Sulfur S2- Sulfide
• Nitrogen N3- Nitride
• Iodine I1- Iodide
• Bromine Br1- Bromide
• Phosphorus P3- Phosphide

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IV. Nomenclature B. Ionic CompoundsExamples

• NaCl
• Na2O
• AlF3
• Be3N2
• Calcium Sulfide
• Barium Iodide
• Barium Oxide
• Magnesium Nitride

• Sodium Chloride
• Sodium Oxide
• Aluminum Fluoride
• Beryllium Nitride
• CaS
• BaI2
• BaO
• Mg3N2

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IV. Nomenclature C. Molecular CompoundsRule 2

• When nonmetals H combine with each other
  through sharing electrons (covalent bonds), they
  form molecules there are no ions.
• Rule 2 When both elements are nonmetals
  (molecular compounds), then Name the the -
  change ending to ide as before. Use prefixes
  of di, tri, tetra, penta, etc to tell how many of
  each element is present.

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IV. Nomenclature C. Molecular Compounds

• CO2 Carbon Dioxide
• CCl4 Carbon Tetrachloride
• N2O Dinitrogen Oxide
• P2S5 Diphosphorus Pentasulfide
• PBr3 Phosporus Tribromide
• BI3 Boron Triiodide
• Notes (1) Organic compounds like CH4 use their
  own rules which we wont cover.
• (2) diatomic molecules named with the
  element name. O2 Oxygen

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V. Polyatomic Ions

• Previous compounds formed from two elements.
• Frequently have compounds formed from three or
  four elements. When this happens, then usually
  have a polyatomic ion present.
• Polyatomic ions stable ions formed from two or
  more elements held together by covalent bonds.
• Examples
• SO4-2 Sulfate NO2- nitrite PO4-3
  Phosphate

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V. Polyatomic Ions

• Polyatomic ions are held together by covalent
  bonds, and they form ionic bonds with metals.
• Examples NaNO2 Na2SO4 Na3PO4
• When have more than one polyatomic ion in a
  compound then use parentheses around the ion.
• Examples Na2SO3 Ca(NO2)2 Ca3(PO4)2
• Nomenclature Simply use the polyatomic ion
  name.
• Example Calcium Nitrite Calcium Phosphate
  above
• Need to memorize the following polyatomic ions,
  their names and their charges.

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V. Polyatomic Ions - Memorize the Names,
Formulas and the Charges

• Formula Name Formula Name
• NH4 Ammonium (The Only Positive One in
  this list)
• C2H3O2- Acetate CN- Cyanide
• NO3- Nitrate NO2- Nitrite
• OH- Hydroxide
• HCO3- Hydrogen Carbonate
• CO3-2 Carbonate Cr2O7-2 Dichromate
• SO4-2 Sulfate SO3-2
  Sulfite
• PO4-3 Phosphate

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V. Polyatomic Ions Examples of Naming and
Obtaining Formulas

• Aluminum Hydroxide
• Calcium Cyanide
• Barium Sulfate
• Ammonium Nitrate
• Ba(OH)2
• LiNO2
• KNO3
• NaHCO3
• Al2(SO4)3

• Al(OH)3
• Ca(CN)2
• BaSO4
• NH4NO3
• Barium Hydroxide
• Lithium Nitrite
• Potassium Nitrate
• Sodium Hydrogen Carbonate
• Aluminum Sulfate

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Naming Mixed Examples

• NaF
• CS2
• NI3
• BaI2
• K3PO4
• Boron Trifluoride
• Sodium Sulfite

• Sodium Fluoride
• Carbon Disulfide
• Nitrogen Triiodide
• Barium Iodide
• Potassium Phosphate
• BF3
• Na2SO3