Chem 20: Unit A

• Bonding results from the simultaneous attraction of electrons by two atomic nuclei
• Ionic bond: electrons are transferred from metal to nonmetal, forming a crystal lattice structure
• Molecular bond: electrons are shared between two nonmetals
• Valence electrons occupy the outermost energy level of an atom
• Electrons exist in orbitals, regions of space around an atom's nucleus where an electron may exist
• Rules of Quantum Mechanics: up to four orbitals can be occupied in the first three periods
• Bonding electron: a single electron occupying an orbital; lone pair: a pair of electrons filling an orbital
• Lewis Dot Diagrams represent the arrangement of electrons in atoms using dots for valence electrons
• Electronegativity is the relative ability of an atom to attract 2 bonding electrons into its valence shell
• Electronegativity increases from left to right in a period and as we move up a group
• Ionic bond forms when two atoms with different electronegativities collide, resulting in the transfer of electrons
• All ionic compounds are arranged in a 3-dimensional crystal lattice structure
• Molecular bonds are present when two atoms have roughly equal electronegativities, while ionic bonds form with distinctly unequal electronegativities
• Octet Rule states that when atoms combine, bonds are formed so each atom (except hydrogen and helium) finishes with 8 electrons
• Double bond involves the sharing of two electron pairs between two atoms; triple bond involves sharing three electron pairs
• The theory of a double bond involves the sharing of two electron pairs between two atoms
• A triple bond involves sharing three electron pairs
• Many atoms can form more than one kind of molecular bond
• The maximum number of single molecular bonds that an atom can form is known as its bonding capacity
• In Lewis structures, lone pairs of electrons are not indicated, while shared pairs of electrons are represented by lines
• To draw Lewis Diagrams (CH3Cl):

  • Add all the valence electrons: C = 4, H = 1 x 3 = 3, Cl = 7, Total = 14 e–
  • Determine the central atom (the least electronegative or most bonding electrons): Carbon
  • Determine total e– for octets: C = 8, H = 2 x 3 = 6, Cl = 8, Total = 22 e–
  • Determine # of bonds needed by dividing bonding e– by 2: 8 e– / 2 = 4 bonds needed
  • Draw single bonds to the central atom, put all remaining valence electrons on atoms as lone pairs, turn lone pairs into double or triple bonds as needed to give every atom an octet
• Empirical evidence indicates that molecules have definite 3-D shapes determined by the Valence Shell Electron Pair Repulsion Theory (VSEPR) developed by R.J. Gillespie
• In VSEPR, the shape of a molecule is determined by the number of pairs of valence electrons, both bonded and unbonded
• Lone pairs of electrons affect the shape of the molecule, for example, in NH3 where there are 3 pairs of bonding electrons, resulting in a trigonal pyramidal shape
• Bond dipoles are used to show if a bond is polar or non-polar, with the arrow pointing from lower to higher electronegativity
• A polar molecule is one where the charge is not distributed equally among the atoms that make up the molecule, resulting from a bond between atoms with unequal electronegativities
• The more uneven the electrons are shared, the more polar the bond will be
• Non-polar molecular bond: a bond between atoms with equal electronegativities where the atoms share their electrons equally
• Molecules with several different polar bonds can still be non-polar if the electronegativities are similar and the atoms are spaced equally around the central atom
• Symmetrical shapes are those without a lone pair on the central atom, like Tetrahedral- CH4, Trigonal planar- BH3, Linear- CO2, which will be nonpolar if all the atoms are the same
• Polar molecules and ions dissolve well in polar solvents, while nonpolar molecules dissolve in nonpolar solvents following the principle "like dissolves like"
• Intra stands for inside/within, while inter is between in the context of intramolecular and intermolecular forces
• Network Covalent Solids are the strongest intramolecular force, composed of a 3D network of covalently bonded atoms like diamond, graphite, and silicon carbide, having high melting points and being often poor conductors of heat and electricity
• Diamond’s structure is an example of a covalent network, where individual carbon atoms are held together by covalent bonds
• Ionic Bonds are composed of a lattice in which each ion is surrounded by neighboring ions with opposite charge, being hard, brittle, having high melting points, and only being conductive in molten liquid or aqueous form
• Polyatomic ions are held together by covalent bonds but form ionic bonds with other ions
• Metallic Solids are composed of a lattice of metal atoms held together by the attraction of the nuclei and the valence electrons, being malleable, ductile, and good conductors like Fe(s) and Al(s)
• Molecular Compounds are the weakest type of intramolecular bonding, made up of individual molecules of covalently bonded atoms, existing as solids, liquids, or gases with relatively low melting points and being poor conductors like H2O(l) and CO2(g)
• London Dispersion Forces are attractive forces caused when electrons in one molecule are temporarily attracted to the protons in another molecule, increasing with the number of protons and electrons in larger molecules
• Dipole–Dipole Forces are the attraction between oppositely charged ends of polar molecules, occurring when polar molecules are attracted to each other and being slightly stronger than dispersion forces