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Covalent Bonding and IMF. Unit 2. 12.1 Types of Chemical Bonds. What is a bond and what holds it together? What are the two types of bonds?. 12.2 Electronegativity. How does electronegativity affect the bond type?. 12.2 Electronegativity. Which elements will form each type of bond?.

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12 1 types of chemical bonds
12.1 Types of Chemical Bonds
  • What is a bond and what holds it together?
  • What are the two types of bonds?
12 2 electronegativity
12.2 Electronegativity
  • How does electronegativity affect the bond type?
12 2 electronegativity1
12.2 Electronegativity
  • Which elements will form each type of bond?
12 3 bond polarity and dipole moments
12.3 Bond Polarity and Dipole Moments
  • When will a dipole moment occur within a compound?
12 3 bond polarity and dipole moments1
12.3 Bond Polarity and Dipole Moments
  • When will binary compound have a dipole moment?
  • When will a polyatomic compound have a dipole moment?
12 6 lewis structures
12.6 Lewis Structures
  • It is a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. The valence electrons are represented by “dots”.
  • Guidelines for drawing:
    • Determine total number of valence electrons that the compound has available for sharing.
    • Bonded atoms will share a pair or more pairs of electrons.
    • Each atom must have a full valence share (OCTET RULE).
      • Exceptions to the octet rule:
12 6 lewis structures1
12.6 Lewis Structures
  • Practice:
    • Binary
    • Polyatomic
  • Lewis Structure Lingo:
    • Lone Pair
    • Bonding Pair
12 7 lewis structures of molecules with multiple bonds
12.7 Lewis Structures of Molecules with Multiple Bonds
  • Single Bond:
  • Double Bond:
  • Triple Bond:
  • Resonance:
  • Other tricky ones: PCl5or XeCl4
12 9 molecular structure vsepr model
12.9 Molecular Structure: VSEPR Model
  • Molecular Structure=Geometric Structure=VSEPR Shape
    • Determining the structure of the molecule is important because its structure plays a part in the molecules chemical properties.
  • VSEPR –
    • Main idea:
  • Steps for Predicting the VSEPR Shape:
  • Practice: BeF2, H2O, BF3, CH4, NH3
12 10 molecular structure molecules with multiple bonds
12.10 Molecular Structure: Molecules with Multiple Bonds
  • When using VSEPR model to predict the molecular geometry of a molecule, a double or triple bond is counted the same as a single electron pair.
  • Practice: [NO3] -1, [SO3]2-
5 3 naming binary compounds that contain only nonmetals
5.3 Naming Binary Compounds that Contain Only Nonmetals
  • The first element is named using the full element name.
  • The second element takes the root name of the element with an “ide” ending.
  • Prefixes are used to denote the number of atoms present in each type.
    • The prefix “mono” is never used on the first element in the formula.
examples and practice
Examples and Practice
  • H2O -Dihydrogenmonoxide
  • CO2 -Carbon dioxide (Do NOT use mono for carbon)
  • P2O5 -DiphosphorusPentaoxide
  • Practice
    • N2O4
    • SCl3
    • SiBr4
    • PCl5
    • SO2
5 7 writing formulas from names
5.7 Writing Formulas from Names
  • Determine the number for each prefix used.
  • Write that number as the subscript for that element.
  • Examples:
    • Diphosphorusdecaoxide is P2O10
    • Oxygen difluoride is OF2
    • Carbon Monoxide is CO
    • Trinitrogenhexaoxide is N3O6
practice
Practice
  • Diphosphorustetraoxide
  • Sulfur dibromide
  • Phosphorus trichloride
  • Dinitrogenoctaoxide
  • Oxygen dichloride
14 3 intermolecular forces
14.3 Intermolecular Forces
  • INTRAmolecular vs. INTERmolecular Forces
  • Intermolecular forces in covalent compounds is the result of shared electrons.
    • Dipole-dipole attraction (polar compounds)
    • Hydrogen bonding (between H and NOF)
    • London dispersion forces (exist in all molecules including nonpolar compounds and noble gases)
slide24
IMF
  • Dipole-Dipole
    • Weaken as distance increases
    • Due to the close proximity of solid molecules, there is a repulsion factor to consider.
slide25
IMF
  • Hydrogen Bonding
    • Hydrogen is bound to highly electronegative atom
    • A particularly strong type of dipole-dipole because:
      • Great polarity of the bond itself
      • Proximity of dipoles due to small size of hydrogen
    • Results in the unusually high boiling point of water
slide26
IMF
  • London Dispersion (weakest IMF)
    • Since all substances exist in the liquid and solid state, we know molecules without dipole moments must exert a force on each other.
    • Electrons are constantly moving and may become unevenly distributed.
      • This creates a temporary dipolar arrangement.
      • This instantaneous dipole can induce a similar dipole in a neighboring atom.
    • These forces become stronger with more electrons.
14 5 the solid state types of solids
14.5 The Solid State: Types of Solids
  • Crystalline solids: regular arrangement of components
    • Molecular solids are comprised of molecules.
      • Polar molecules held together by H bonding or dipole-dipole AND London Dispersion
      • Nonpolar molecules held together by London Dispersion
    • Atomic solids are comprised of one element covalently bound to each other.
      • Diamonds are covalently bound carbon atoms.
14 6 bonding in solids
14.6 Bonding in Solids
  • Molecular solids:
    • The weaker the force, the easier that the solid will melt (they generally have low melting points).
    • If dissolved in water, the solid splits into individual molecules and will not conduct electricity.
  • Atomic solids:
    • Properties depend on the atoms themselves.
    • Metals have a “sea of electrons” that determine their properties.
    • Allotropes of Carbon
20 1 carbon bonding
20.1 Carbon Bonding
  • Silicon (geological world):
  • Carbon (biological world):
20 2 alkanes
20.2 Alkanes
  • Saturated Hydrocarbons
  • VSEPR Shape:
  • General Formula:
prefixes1
Prefixes
  • Methane (1)
  • Ethane (2)
  • Propane (3)
  • Butane (4)
  • Pentane (5)
  • Hexane (6)
  • Heptane (7)
  • Octane (8)
  • Nonane (9)
  • Decane (10)
20 4 naming alkanes
20.4 Naming Alkanes
  • For nomenclature rules see pg. 616
    • Identify LONGEST chain of carbons
    • Name that chain.
    • Circle all alkyl groups and halogens attached to the chain
    • Number the chain
    • Name these functional groups
20 7 alkenes and alkynes
20.7 Alkenes and Alkynes

Alkenes (CnH2n)

  • VSEPR Shape
slide35
Nomenclature is similar to alkanes except:
    • Root changes from “ane” to “ene”.
    • Location of the double bond is indicated by the lowest-numbered carbon involving the bond. DOUBLE BOND MOST IMPORTANT!
      • 1-butene (not 3 butene)
      • 2-butene
20 7 alkenes and alkynes1
20.7 Alkenes and Alkynes
  • Alkynes
    • VSEPR Shape:
    • General Formula:
    • Nomenclature is similar to alkanes except:
      • Root changes from “ane” to “yne”
      • Location of the triple bond is indicated by the lowest-numbered carbon involving the bond.
20 8 aromatic hydrocarbons
20.8 Aromatic Hydrocarbons
  • C6H6 Benzene
    • Resonance structures
    • Six delocalized electrons
substituted benzene rings
Substituted Benzene Rings
  • Ortho (o)
    • 1, 2 substituted
  • Meta (m)
    • 1, 3 substituted
  • Para (p)
    • 1, 4 substituted
20 10 functional groups
20.10 Functional Groups
  • Alcohols (R-OH)
  • Ethers (R-O-R’)
  • Aldehydes (R-COH)
  • Ketones (R-CO-R’)
  • Carboxylic acids (R-COOH)
  • Esters (R-COO-R’)
  • Amines (R-NH2)
common polymers
Common Polymers
  • Definition
  • Common Types:
    • Proteins
    • Nucleic Acids
    • Carbohydrates
    • Plastics