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Covalent Bonding and IMF

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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Covalent Bonding and IMF

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  1. Covalent Bonding and IMF Unit 2

  2. 12.1 Types of Chemical Bonds • What is a bond and what holds it together? • What are the two types of bonds?

  3. 12.2 Electronegativity • How does electronegativity affect the bond type?

  4. 12.2 Electronegativity • Which elements will form each type of bond?

  5. 12.3 Bond Polarity and Dipole Moments • When will a dipole moment occur within a compound?

  6. 12.3 Bond Polarity and Dipole Moments • When will binary compound have a dipole moment? • When will a polyatomic compound have a dipole moment?

  7. 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:

  8. 12.6 Lewis Structures • Practice: • Binary • Polyatomic • Lewis Structure Lingo: • Lone Pair • Bonding Pair

  9. 12.7 Lewis Structures of Molecules with Multiple Bonds • Single Bond: • Double Bond: • Triple Bond: • Resonance: • Other tricky ones: PCl5or XeCl4

  10. 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

  11. A linear molecule, e.g. BeF2.

  12. A bent or V-shaped molecule, e.g. H2O.

  13. A trigonal planar (triangular) molecule, e.g. BF3.

  14. A tetrahedral molecule, e.g. CH4.

  15. A trigonal pyramid, e.g. NH3.

  16. 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-

  17. 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.

  18. Prefixes

  19. Examples and Practice • H2O -Dihydrogenmonoxide • CO2 -Carbon dioxide (Do NOT use mono for carbon) • P2O5 -DiphosphorusPentaoxide • Practice • N2O4 • SCl3 • SiBr4 • PCl5 • SO2

  20. 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

  21. Practice • Diphosphorustetraoxide • Sulfur dibromide • Phosphorus trichloride • Dinitrogenoctaoxide • Oxygen dichloride

  22. 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)

  23. IMF • Dipole-Dipole • Weaken as distance increases • Due to the close proximity of solid molecules, there is a repulsion factor to consider.

  24. 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

  25. 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.

  26. 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.

  27. 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

  28. 20.1 Carbon Bonding • Silicon (geological world): • Carbon (biological world):

  29. 20.2 Alkanes • Saturated Hydrocarbons • VSEPR Shape: • General Formula:

  30. Prefixes • Methane (1) • Ethane (2) • Propane (3) • Butane (4) • Pentane (5) • Hexane (6) • Heptane (7) • Octane (8) • Nonane (9) • Decane (10)

  31. Properties of Alkanes

  32. 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

  33. 20.7 Alkenes and Alkynes Alkenes (CnH2n) • VSEPR Shape

  34. 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

  35. 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.

  36. 20.8 Aromatic Hydrocarbons • C6H6 Benzene • Resonance structures • Six delocalized electrons

  37. Substituted Benzene Rings • Ortho (o) • 1, 2 substituted • Meta (m) • 1, 3 substituted • Para (p) • 1, 4 substituted

  38. 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)

  39. Common Polymers • Definition • Common Types: • Proteins • Nucleic Acids • Carbohydrates • Plastics

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