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Chapters 4 & 5

Chapters 4 & 5. Chemical Bonding. Valence Electrons. Outermost electrons s and p electrons for main group elements Responsible for chemical properties of atoms Participate in chemical reactions. Valence Electron. Core Electrons. Problems.

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Chapters 4 & 5

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  1. Chapters 4 & 5 Chemical Bonding

  2. Valence Electrons • Outermost electrons • s and p electrons for main group elements • Responsible for chemical properties of atoms • Participate in chemical reactions Valence Electron Core Electrons

  3. Problems • Write out the electron configurations for the following elements and identify how many core and valence electrons each has. • Mg • S • Br • Kr

  4. Lewis Dot Structures • LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

  5. Problems • Write the Lewis Dot Structures for the following atoms • Li • Be • Br • C • N • Ne

  6. Li: [He]1s1 Na: [Ne]2s1 K: [Ar]3s1

  7. Octet Rule • Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells • Natural electron configuration of the Noble Gases • Done by gaining, losing, or sharing electrons • Increases stability • H and He seek a “Duet”

  8. Ionic Bonding • Ions: atoms that have a charge due to gain or loss of electrons • Anion: (-) charged atom • Cation: (+) charged atom • Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

  9. Formula Unit

  10. Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other • Metal + Non-metal • NaCl • Metal + Polyatomic Ion • NaNO3 • Polyatomic Ion + Non-metal • NH4Cl • Polyatomic Ion + Polyatomic Ion • NH4NO3 • Net charge on compound equal to zero

  11. Oxyanions

  12. Rules For Naming Ionic Compounds • Name the cation by its elemental/polyatomic name • If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses • Next, name the anion and change its ending to “-ide” • If the anion is polyatomic, do not change the ending to “-ide” • Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

  13. Problems Write the name for the following compounds: • KI • MgBr2 • Al2O3 • FeCl2 • CaSO4 • Ba(NO2)2 • Cu(NO3)2

  14. Write the Formula for the following ionic compounds: • Sodium Fluoride • Calcium Sulfite • Calcium Chloride • Iron (III) Oxide • Cobalt (II) Hydroxide • Ammonium Bromide • Ammonium Carbonate • Aluminum Carbonate

  15. Iron (II) Chloride Iron (III) Chloride

  16. Covalent Compounds • Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons • Electrons NOT transferred • No + or – charges on atoms • Non-metal + Non-metal • Also called “molecules” • Examples: • H2O • CO2 • Cl2 • CH4

  17. orH-H Duet or

  18. Name the first non-metal by its elemental name Add a prefix to indicate how many Name the 2nd non-metal and change its ending to “-ide” Add a prefix to indicate how many Naming Covalent Compounds

  19. Problems Write the name of the following compounds: • CO • NI3 • N2O • SF6 • B2O3

  20. Write the formula for the following compounds: • Phosphorous Pentachloride • Nitrogen Monoxide • Dinitrogen Tetroxide • TetraphosphorousDecoxide

  21. Problems • KCl • Na2S • H2O • SO2 • K3PO4 • FeCl3 • (NH4)2SO4 • SCl2 • Cu(OH)2 • P2O5

  22. Sodium Iodide • Aluminum Sulfate • Phosphorous Pentabromide • Magnesium Nitride

  23. Naming Acids • Acids that do not contain oxygen • Begin the name with “hydro” • Name the anion, but change the ending to “-ic” • Add “acid” on the end • HCl • HF

  24. Acids that contain oxygen • Do not put “hydro” at the beginning • Begin the name with the anion • If the anion has the ending “-ate,” change this to “-ic acid” • If the anion has the ending “-ite,” change this to “-ous acid” • HClO4 • HClO3 • HClO2 • HClO

  25. Problems • Name the following • HBr(g) • HBr(aq) • HNO2(aq) • HNO3(aq) • HI (aq) • HI (g) • H2CO3 (aq) • H3PO4 (aq) • H3PO3 (aq) • HCN (aq)

  26. Molecular Structures

  27. Ball & Stick Models Space-Filling Models Water Methane

  28. Ethanol

  29. Lewis Dot Structures • Count the total number of valence electrons in the molecule. Ex: PCl3 • Use atomic symbols to draw a proposed structure with shared pairs of electrons. • Atoms don’t tend to bond to other atoms of the same element when they can avoid it • Exception: Carbon

  30. Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms • Place any leftover electrons on the central atom • If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple). • Ex: CH2O

  31. Problems Draw the LDS’s for the following molecules: • Cl2O • C2H4 • C2H6O

  32. What Things Like To Do • Halogens • Like to be terminal • Like to have one single bond and 3 lone pairs (non-bonding electrons) • Carbon • Likes to have 4 single bonds and no lone pairs • A double bond counts as two singles • A triple bond counts as three singles • Likes to be central • Likes to bond to other carbons

  33. Silicon • Likes to do what carbon does • Oxygen • Likes to have two single bonds and 2 lone pairs • Sulfur • Likes to do what oxygen does • Nitrogen • Likes to have 3 single bonds and one lone pair

  34. Phosphorous • Likes to do what nitrogen does • Hydrogen • Likes to be terminal with only one single bond • No lone pairs!

  35. Problems • SH2 • C3H8 • Si2H6 • PI3 • CH3OH • C2H2

  36. CCl2O • N2H4 • CH2OS • C2H6O • CO • BrHO

  37. Electronegativity • The measure of the ability of an atom to attract electrons to itself • Increases across period (left to right) and • Decreases down group (top to bottom) • fluorine is the most electronegative element • francium is the least electronegative element

  38. Electronegativity Scale

  39. ENCl = 3.0 3.0 - 3.0 = 0 Pure Covalent Types of Bonding • Non-Polar Covalent Bond: • Difference in electronegativity values of atoms is 0.0 – 0.4 • Electrons in molecule are equally shared • Examples: Cl2, H2, CH4

  40. ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent • Polar Covalent Bond: • Difference in electronegativity values of atoms is 0.4 – 2.0 • Electrons in the molecule are not equally shared • The atom with the higher EN value pulls the electron cloud towards itself • Partial charges • Examples: HCl, ClF, NO

  41. ENCl = 3.0 ENNa = 1.0 3.0 – 0.9 = 2.1 Ionic • Ionic Bond: • Difference in EN above 2.0 • Complete transfer of electron(s) • Whole charges

  42. Problems • Predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable • KBr • HF • BrI • FI

  43. Valence Shell Electron Pair Repulsion Theory • VSEPR theory: • Electrons repel each other • Electrons arrange in a molecule themselves so as to be as far apart as possible • Minimize repulsion • Determines molecular geometry

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