Unit A: Bonding 3.1-3.2

• Frankland


  each element has a fixed bonding capacity
• Kekule


  illustrated a bond as a dash between bonding atoms
• Hoff and Le Bel


  -extended diagrams to 3D

  -revised a theory to explain the ability of certain substances to change light as it passes through a sample of the substance
• Abegg


  -suggested that bonding capacity must somehow be associated with an atoms electron structure.

  -stability of the noble gases was due to the number of electrons in the atom.
• Lewis


  -proposed that atoms could achieve stable electron arrangements by: sharing electrons as well as by transferring them.
• Ionic Bond


  simultaneous attraction between positive and negative ions.
• Covalent Bond


  simultaneous attraction of the nuclei of two atoms for valence electrons that they share between them.
• Quantum Mechanics


  -mathematical model

  -electrons are described in terms of their energy content, and orbitals in terms of calculated probability of an electron being at any given point relative to the atomic nucleus.
• Pauling


  -explained why certain electron arrangements are stable.

  -showed that electron sharing must cover a complete range from equal attraction to total transfer.
• Orbital 


  a specific volume of space in which an electron of certain energy is likely to be found

  contain: 2, 1 or 0 electrons
• Valence Orbital


  Space that can be occupied by electrons in an atoms highest energy level.
• Bonding Electron


  Full valence orbital occupied by 3 electrons.
• Lone Pair


  Two electrons occupying the same orbital.
• electrons are not dots and are not stationary
• the farther away from the nucleus that electrons are, the weaker their attraction to the nucleus
• Inner electrons shield the valence electrons from the attraction of the positive nucleus.
• The greater the number of protons in the nucleus, the greater the attraction for electrons must be.
• Electronegativity


  the relative ability of an atom to attract a pair of bonding electrons in its valence level.
• Metals tend to have low electronegativities
• Non-metals tend to have high electronegativities.
• Covalent Bonding


  -attraction of two nuclei for a shared pair of bonding electrons.

  -forms between two non metal atoms

  -products are molecular substances
• If the electronegativities of both atoms are relatively high, neither atom will "win" and the pair of bonding electrons will be shared between the two atoms.
• Ionic Bond


  attraction between any specific cation and any specific anion.
• after electron transfer ions arrange so that maximum total attraction between positive and negative charges occur.
• Ions Goal


  have a total net charge of zero.
• Metallic Bonding


  If both types of colliding atoms have relatively low electronegativities, the atoms can share valence electrons. No chemical reaction occurs.
• In metallic bonding, the valence electrons are not held very strongly by their atoms. So valence electrons can move freely between the positive ions and the negative ions.
• wherever the electrons move they are acting to hold atoms together because positive nuclei on either side will both attract the electrons.
• Bonding in Metallic substances


  -great number of positive ions surrounded by a "sea" of mobile electrons. Valence electrons act like glue (hold s all together)
• Attraction force around a metal atom acts in every direction.
• Bonding Theory 


  That a covalent bond between the atoms results from the simultaneous attraction of two nuclei for a shared pair of electrons.
• Double Covalent Bond


  sharing two pairs of electrons at once to from a double covalent bond.
• Triple Bond


  Two atoms sharing three pairs of electrons.
• Double and triple bond is only 1 bond. Lines tell us how many electrons are shared by the bonded atoms
• Molecular Compounds


  Cannot be represented by a simplest ratio formula.
• Simplest Ratio Formula


  Indicate only the relative numbers of atoms or ions in a compound.
• Molecular formulas also accurately represent the actual composition of the smallest units of molecular compounds.
• Bonding Capacity


  maximum number of single covalent bonds that an atom can form.
• Coordinate Covalent Bond


  covalent bond in which one of the atoms donates both electrons.
• Molar mass of an ionic compound is the mass of a mole of formula units.