CHEMSITRY - CHAPTER #6 Chemical Bonding Mr. Novak’s Chapter #6 NOTES/OUTLINE

  1. Groups - columns going down the periodic table. Elements within a group have similar properties because they have the same number of valence electrons.
  2. Valence Electrons are in the highest occupied energy level (outer shell). They determine the physical and chemical properties of the element.
  3. The group number for the representative element tells the number of valence electrons.

Group 1A has one valence electron.

Group 2A has two valence electrons.

Group 3A has three valence electrons.

Group 4A has four valence electrons.

Group 5A has five valence electrons.

Group 6A has six valence electrons.

Group 7A has seven valence electrons.

Group 8A has eight valence electrons (except Helium with only 2).

  1. Lewis Electron Dots show the valence electrons of the atom.
  2. Inner electrons and the nuclei are represented by the element’s symbol.
  3. Bots representing electrons are arranged symmetrically around symbol.
  4. Bonds are formed between atoms using unpaired valence elecctrons
  5. Write the electron dot structures for the following compounds

Water H2O

Ammonia NH3

Methane CH4

Carbon Dioxide CO2

  1. Noble Gases are inert – they are extremely stable and do not react under normal laboratory conditions. They have eight electrons in their highest energy level – eight valence electrons is a stable electron configuration.
  2. Octet Rule - atoms in compounds tend to create the electron configuration of a noble gas –eight electrons in their highest occupied energy level.
  3. Metals lose electrons to obey this rule and become stable like a noble gas.
  4. Nonmetals gain or share electrons to obey the octet rule, become stable.
  5. Cations - neutral atom lost valence electronsbecame a positively charged ion.
  6. Metalswith up to three valence electrons that can be easily removed. (Group 1A, 2A, 3A)
  7. Unstable neutral atoms lose enough electrons to fulfill octet rule in the next energy level down to become more stable as positively charged ions.
  8. Proton number never changes. A neutral atom has equal numbers of protons and electrons. If electrons are lost then there are less electrons than protons and the atom is now positively charged and called a cation.
  9. Examples: K loses 1e- to attain the stable noble gas configuration of Argon. Al loses 3e- to attain the stable noble gas configuration of Neon. (Board Game Analogy)
  10. Anions - neutral atom gained valence electrons& became a negatively charge ion.
  11. Nonmetalswith up to three valence electrons that can be easily added. (Group 5A, 6A, 7A)
  12. Unstable neutral atoms gain enough electrons to fulfill octet rule in their highest energy level and become more stable as negatively charged ions.
  13. Proton number never changes. A neutral atom has equal numbers of protons and electrons. If electrons are gained then there are more electrons than protons and the atom is now negatively charged and called an anion.
  14. Examples: Cl gains 1e- to attain the stable noble gas configuration of Argon. N gains 3e- to attain the stable noble gas configuration of Neon. (Board Game Analogy)
  15. Polyatomic Ions – a group of atoms that acts as a unit with a single charge
  16. Begin memorizing polyatomic ions…get the list from the website and make flashcards.
  17. Know the formula, the charge, and the correct spelling of the name of the seventeen polyatomic ions listed on the website.
  18. Ionic Compound – between metal (cation) and a nonmetal (anion) = (M)(NM)
  19. Ionic Bond - Completely transfer electrons.
  20. Positive charge – cation – lost electrons to the anion.
  21. Negative charge – anion – gained electronsfrom the cation.
  22. Positive charge must equal and, therefore, cancel the negative charge. Example: Sodium Chloride – sodium wants to lose one electron to become stable and chlorine wants to gain one electron to become stable. Will fulfill the octet rule once they combine. (+)(-) = 0
  23. Formula unit – a chemical formula of the smallest sample of an ionic compound.
  24. Properties of Ionic Compounds
  25. Crystalline solids at room temperature. Arranged in repeating three-dimensional patterns. Very stable. Structures determined by X-ray diffraction crystallography. Example: In solid NaCl, each Na is surrounded by six Cl and each Cl is surrounded by six Na.
  26. Have very high melting points. Separates each ion from one another. Hard to break the attraction between the ions. Very stable. Example: NaCl melts at 800 ˚Celsius.
  27. Conduct electric currents when molten (liquid) or dissolved in water (aqueous). The cations and anions migrate freely.
  28. Ionic compounds – are electrically neutral salts. (Many appear as minerals in the Earth’s crust.)
  29. Ionic Character
  30. Ionic compounds have the greatest ionic character with full on charged ions. The further the ions are apart in electronegativity, the more the ionic character.
  31. Molecular compounds have very low electronegativity. The closer the ions are in electronegativity, the less the ionic character.
  32. Molecular Compounds - (NM)(NM) – are often are multiples of the lowest whole-number ratios of nonmetals. Examples: C3H6 and C4H10 Note: Do not reduce molecular compounds.
  33. Covalent Bonds - the sharing of electrons between two nonmetals – creates a molecular compound (or molecule). The goal is to attain eight valence electrons –stability – similar to a noble gas electron configuration.
  34. Do not forget that Hydrogen is a nonmetal.
  35. Do not forget your diatomic molecules in Group 7A: N2, O2, F2, Cl2, Br2, I2, and H2
  36. Lewis Structures
  37. Shared Pairs – both atoms can claim the electrons to achieve the octet rule and become like noble gas configurations.
  38. Unshared Pairs - pairs of valence electrons that are not shared between atoms – also called lone pairs. Example: F2
  39. Single Covalent Bond - formed when one pair of electrons is shared between two atoms.Example:Hydrogen – diatomic molecule…H2
  40. Double Bond - involve two shared pairs of electrons. (Oxygen will form the double bond but be an exception to the octet rule.) Example: Oxygen - diatomic molecule…O2
  41. Triple Bond -involve three shared pairs of electrons. Example: Nitrogen – diatomic molecule…N2
  42. Structural Formulas - chemical formulas that show the arrangement of atoms in molecules. A dash represents a pair of shared electrons (never used to show ionic bonds because ions do not share electrons).
  43. No unshared (lone) pairs visible.
  44. Shared pairs: represented with a line (dash) instead of two dots.
  45. Resonance- when two or more electron dot structures can be written for a molecular compound. Example: NO2
  46. Exceptions to the Octet Rule
  47. It is impossible to fulfill the octet rule whenever the total number of valence electrons in the compound is an odd number. Oxygen and Boron tend to be satisfied with less than eight valence electrons. (6-7) Example: BF3
  1. Phosphorus and Sulfur tend to accept more than eight valence electrons because they have the d sublevel to expand into. (10-12) Example: SF6 and PCl5
  2. Properties of Molecular Compounds
  3. Do not conduct electricity.
  4. Solids, Liquids, and gases.
  5. Low melting and boiling points that only separate one molecule from another as opposed to separating each atom from another.
  6. VSEPR Theory - valence shell electron pair repulsion. The electron dot structures are not flat 2D structures, but are 3D in real life.
  7. Shapes - Electron pairs repel. Molecules adjust their shapes so that the valence electron pairs are as far apart as possible.
  8. Linear – angles are 180 degrees – definitely will be linear if only have two atoms in the molecule. No lone pairs and two covalent bonds or three lone pairs and one covalent bond around central atom. Example: CO2
  9. Bent – again, unshared pair(s) strongly repels the covalent bonding pairs. Two lone pairs and two shared pairs around central atom. All angles are 105 degrees. Example: H2O
  10. Trigonal-Planar – three shared pairs(covalent bonds ) separate as much as possible, but are unaffected by a lone pair(no lone pairs) of electrons like the pyramidal structure. Example: BF3
  11. Trigonal-Pyramidal – one unshared pair strongly repels the three shared pairs (covalent bonding), pushing them closer together. All angles are 107 degrees. Example: NH3
  12. Tetrahedral – four faced – four shared pairs and no lone pairs, all angles are 109.5 degrees. Example: CH4
  13. TrigonalBipyramidal – five shared pairs separate as much as possible, but are unaffected by a lone pair of electrons(no lone pairs). Example: PCl5
  14. Octahedral– six shared pairs separate as much as possible, but are unaffected by a lone pair of electrons (no lone pairs). Example: SF6
  15. Can predict the shape of the molecule (as a general rule) using the group number (valence electrons)

1A LinearExample: Li

2A LinearExample:Be

3A BentExample:B

4A TetrahedralExample:C

5A PyramidalExample:N

6A BentExample:O

7A BentExample:F

  1. Hybridization - two atoms combine, their atomic orbitals overlap to produce molecular orbitals. One electron from each atomic orbital combines to create a shared pair in a molecular orbital.
  2. sp3 hybridization – has electrons in 4 orbitals
  3. sp2 hybridization – has electrons in 3 orbitals
  4. sp hybridization – has electrons in 2 orbitals
  1. Polarity – nonmetals of unequal strength (electronegativity) do not share electrons equally. Note: This only applies to Molecules.
  2. Polar Molecules - has a polar bonds, one end of the molecule has a slightly negative charge while the other end has a slightly positive charge. Dipole – a molecule with oppositely charged ends.
  3. Polar Covalent Bonds - when the atoms are of different types, the bonding electrons are shared unequally. The atom with the stronger electronegativity acquires a slightly negative charge as it draws the electrons toward itself. The atom with the lower electronegativity acquires a slightly positive charge as the electrons are drawn away from it. (Reminder: Electronegativity is the ability of the atom to attract electrons to itself.)
  4. Examples: HCl has 2 different types of atoms. Chlorine is more electronegative than Hydrogen. Chlorine pulls the shared pair closer to its own nucleus creating a partial negatively charge pole. Hydrogen allows the shared pair to be pulled farther from its own nucleus creating a partially positively charge pole. H2O has 2 different types of atoms. Oxygen is more electronegative than Hydrogen. The bent shape due to the lone pairs creates oppositely charged ends.
  5. The greater the difference in electronegativity the greater the polarity.
  6. Example: SO2 is less polar than H2O. Sulfur and Oxygen are located close together on PT small difference in electronegativity slightly polar molecule. Hydrogen and Oxygen are located far apart on PT large difference in electronegativity highly polar molecule.
  7. Polar compounds are high in ionic character due to the partially charged poles on the molecule.
  8. The greater the difference in electronegativity, the greater the polarity on the molecule and the more ionic character the molecule has.
  9. Nonpolar Molecules – Either the molecule has no oppositely charged ends or the ends cancel each other out.
  10. Polar Covalent Bonds cancel each other out.Nonpolar molecules may have polar bonds but the overall molecule is nonpolar because polar ends cancel. Note: It will be a nonpolar molecule if the molecule is symmetrical in 3D and all bonds are exactly the same.Example: CO2
  11. Nonpolar Covalent Bonds - when atoms are the same type, they share the bonding electrons equally. Because have same electronegativity. This is the case with all the diatomic molecules.H2 O2 N2 etc.
  12. Nonpolar compounds have very little ionic character as they sometimes exhibit dispersion forces when their ions vibrate to create a momentary dipole.
  13. van der Waals Forces – intermolecular attractions - attractions between molecules – weak compared to ionic or covalent bonds but still substantial in strength
  14. Dipole Interactions - when polar molecules are attracted to one another: opposite charged regions of polar molecules are attracted.
  15. Hydrogen Bonds – a particularly strong dipole interaction specifically involving hydrogen at the partially positive pole.Hydrogen is covalently bonded to a very electronegative atom AND to an unshared pair of another atom.
  16. Hydrogen is able to bond with the unshared pair of electrons from another molecule because its’ valence electrons are not shielded from the nucleus by another layer of electrons (hydrogen’s valence electrons are directly up against the nucleus). Example: H2O
  17. The more electronegative the element that hydrogen is bonded to the stronger the intermolecular attractions. Example: Which element has stronger intermolecular interactions: H2S or H2O?
  18. Dispersion Forces - weakest of all molecular interactions – caused by the motion of electrons.
  19. Vibrating electrons may end up moving randomly closer to one atom or another creating a momentary dipole.
  20. The more electrons the greater the interaction between nonpolar molecules.

Examples: Halogens

F2 – molecules not touching - gas

Br2 – molecules touching and sliding - liquid

I2 – touching and vibrating in place - solid

  1. Metallic Bondinggroups of closely packed cations in a “sea” of free moving valence electrons
  2. Valence Electrons create an attraction between the free moving valence electrons in the positively charged metal cations
  3. Properties of Metals
  4. Good conductors of electricity – electrons enter one end of the metal bar and leave the other.
  5. Ductile – can be stretched into wires.
  6. Malleable– can be pounded into shapes. Metals ions slide passed one another in a sea of drifting