Chemistry Bonding HL

  1. What is an ionic bond?
    An Ionic Bond is the electrostatic attraction experienced between the electrical charges of a cation and an anion. They are formed when one or more electrons are transferred from one atom to another. The driving force for this transfer is usually the formation of a noble gas configuration. Ionic compounds are typically solids with lattice-type structures. (kup)
  2. What is coordination number?
    Coordination number is used to express the number of ions that surround a given ion in a lattice
  3. What is lattice energy?
    Lattice energy is a measure of the strength of attraction between the ions within the lattice. This is greater for smaller and higher charged atoms.
  4. Explain the melting points of ionic compounds?
    Ionic compounds have high melting points due to strong electrostatic forces of attraction between the ions in their lattice structures.(is directly proportional to interacting charges, so mg 2 and o 2 have a higher melting point than na 1 cl 1)
  5. What is volatility? What is the volatility of ionic compounds generally like?
    Volatility is the tendency for a substance to vaporize. It's low for ionic compounds.
  6. How do we determine whether a non-metal x non-metal compound is ionic?
    Electronegativity difference must be more than 1.8 in order to be ionic
  7. Explain the solubility of ionic compounds.
    Ionic compounds are very soluble, as their ions get hydrated (O is negative, H is positive)
  8. Explain the conductivity of ionic compounds.
    In their solid form, electrons are fixed so they do not conduct electricity. However, in their molten or solution state, the electrons are free to move and conduct electricity.
  9. What is a covalent bond?
    A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei. (a region of relatively high electron density between nuclei that arises at least partly from shared electrons)
  10. How do bonds change in covalent bonds?
    Bond length decreases and bond strength increases as the number of shared electrons increases.
  11. What is the condition for pure covalent compounds?
    If the difference in electronegativity is 0 it’s pure covalent.
  12. An important note about lewis structures is that:
    When you write Lewis structures of cations or anions you should always include square brackets with their charge and their shit. Subtract the ion from the total number of electrons.
  13. VSEPR Theory
    • LP|LP > LP|BP > BP|BP
    • What is the effect of molecular size on melting point?
    • Increasing molecular size will increase melting point and stuff.
  14. Charge of Pb is
  15. Charge of Ag is
  16. What is Phosphate
    PO-3 4
  17. Which atoms can expland their octets?
    Anything other than C N O F can expand
  18. 2 Regions of Electron Density:
    Linear with 180 degrees
  19. 3 Regions of Electron Density 3 bonding 0 lone
    trigonal planar (120 degrees)
  20. 3 Regions of Electron Density 2 bonding 1 lone
    • bent
    • less than 120 degrees
  21. 4 Regions of Electron Density 4 bonding 0 lone
    tetrahedral 109.5 degrees
  22. 4 Regions of Electron Density 3 bonding 1 lone
    pyramidal (107 degrees)
  23. 4 Regions of Electron Density 2 bonding 2 lone
    bent (105 degrees)
  24. 5 Regions of Electron Density 5 bonding 0 lone
    • trigonal bipyramidal
    • 90 degrees and 120 degrees
  25. 5 Regions of Electron Density 4 bonding 1 lone
    • Seasaw
    • 90 degrees and 118 degrees between the two bonded in the middle
    • it's like bend but has two sticking out and the inbetween is like 118 degrees
  26. 5 Regions of Electron Density 3 bonding 2 lone
    • T shaped
    • 90 degrees
  27. 5 Regions of Electron Density 2 bonding 3 lone
    • linear
    • 180 degrees
  28. 6 Regions of Electron Density 6 bonding 0 lone
    • octahedral
    • 90 degrees
  29. 6 Regions of Electron Density 5 bonding 1 lone
    • it loses one of the things sticking out from down under and becomes a square pyramidal
    • with 90 degrees
  30. 6 Regions of Electron Density 4 bonding 2 lone
    • it loses both things sticking out from down and up and turns into
    • square planar
    • with about 90 degrees
  31. What is Nitride
    just N
  32. Sn ion
    +2 +4
  33. ammonium
    NH4 +1
  34. When do electrons become delocalized?
    Electrons can become delocalized when there is more than one possible position for a double bond. These structures are known as resonance structures.
  35. What are giant molecular or network covalent molecules?
    Some covalent structures make giant molecular or network covalent or macromolecular structures.
  36. What are allotropes
    Allotropes are different forms of an element in the same physical state.
  37. Carbon has 4 allotropes:
    • Graphite
    • Diamond
    • Fullerene
    • Graphene
  38. Graphite:
    • Each atom is sp2 hybridized
    • made up of layers held together by london forces so they slide off each other easily
    • good electricity conductor due to one delocalized electron per molecule
    • not a good heat conductor, tends to conduct in a parallel manner
    • most stable allotrope of carbon
  39. Diamond:
    • Each atom is sp3 hybridized
    • tetrahedrally arranged
    • doesn’t conduct electricity
    • very efficient thermal conductor
    • hardest thing ever
  40. Fullerene:
    • sp2 hybridized fixed formula in a sphere
    • low thermal and electrical conductivity
    • reacts with K to make superconducting crystalline material
  41. Graphene:
    • sp2 but no layers, single layer
    • super conducting electricity
    • best thermal conductivity
    • thinnest material ever to exist
  42. silicon also has a giant covalent structure
    SiO2 (glass) is just O in between the tetrahedral Si structure
  43. Why doesn’t diamond conduct electricity as well as other allotropes?
    Diamond doesn’t conduct electricity as well as graphene or graphite bc it has no delocalized electron.
  44. What are Intermolecular Forces:
    • London Forces: temporary, instantaneous dipole, induced dipole (stronger as number of electrons increases)
    • Dipole-dipole: permanent polar
    • both are van der waals
    • 3. Hydrogen Bonding: H must be bonded to FNO (strongest) Boiling point increases as molecular mass increases, however hydrogen bonds break these rules
  45. Strength of intermolecular forces from strongest to weakest?
    hydrogen bonds> dipol-dipol > london forces
  46. How does polarity affect solubility?
    Polar is soluble in polar, non-polar in non-polar however larger molecules with small polar sections are less soluble. Some polar compounds ionize in water and are able to conduct electricity
  47. What is the relationship between volatility and volatility?
    non-polar compounds are the most volatile
  48. What is Metallic bonding:
    Electrostatic attraction between the lattice of cations and delocalized electrons
  49. The strength of metallic bonding is determined by:
    • number of delocalized electrons (more the better)
    • the charge on the cation (bigger the better)
    • the radius of the cation (smaller the better)
  50. Why do transition metals have strong metallic bonds?
    Transition metals have very strong metallic bonds due to many delocalized electrons
  51. What are alloys? How are they made? What are their properties?
    Alloys are made by adding on metal to another metal. Alloys are more chemically stable, durable and resistant to corrosion due to the different packing of cations in the alloy.
  52. How is an expanded octet possible?
    Expanded octet: possible with atoms after the third period bc the d orbital is relatively close to the p orbital
  53. What is formal charge?
    • Number of electrons missing on the atom in the lewis structure
    • Lowest formal charge and
    • negative values on the more electronegative atoms
  54. What are Different kinds of covalent bonds:
    Sigma bonds and pi bonds
  55. Explain Sigma bonds:
    All single bonds are sigma bonds. Sigma bonds form by overlap of orbitals along the bond axis. Electron density is concentrated between the nuclei of the bonded atoms.
  56. Explain Pi bonds:
    When two orbitals overlap sideways, the electron density of the molecular orbital is concentrated in two regions, above and below the plane of the bond axis. Only in double and triple bonds, but they also must include a sigma bond.
  57. Hybridization:
    The combination of two or more orbitals to form orbitals with lower energy. The formation of covalent bonds starts with excitation of electrons to from s orbitals to p orbitals.
  58. SP3
    4 electron domains Sp3: When carbon forms 4 single bonds: tetrahedron 109.5 degrees
  59. SP2
    3 electron domains Sp2: 120 degress Carbon forms a double bond 1 unhybridized p orbital 3 hybridized sp orbitals
  60. SP
    2 electron domains Sp: triple bond two sp orbitals two unhybridized orbitals 180 degrees
  61. How to determine wavelength required?
    h x c x mol / E
Card Set
Chemistry Bonding HL