Carbon may form a single, double and triple bonds to another carbon:
2. Structural or Constitutional Isomers: compounds with the same molecular formula, but different atom connectivity.
Two different compounds with different physical & chemical properties: Ethyl alcohol (ethanol): B.P. 78.5 oC (liquid) M.P. –117.3 oC Reacts with alkaline metals. Dimethyl ether: -24.9 oC (gas!) -138 oC Does not
How many isomers can you draw for:
3. Chemical Bonds Ionic: forms by a complete transfer of one or more electrons from one atom to another to form ions (cation and anion). Forms between typical metal and non-metal due to a significant difference in electronegativity (ability of atom to attract electrons; refer to table 1.2).
Covalent: forms by sharing electrons between two atoms (polar covalent and non-polar covalent). Forms between non-metals. Electronegativity difference ∆ EN for two elements: ∆ EN > 1.2 Ionic Bond 0.4 < ∆ EN < 1.2 Covalent Polar ∆ EN < 0.4 Covalent Non-polar
Octet Rule: to form compounds atoms gain, lose or share electrons in order to achieve a stable electron configuration of 8 electrons in the outer shell. It is only 2 electrons for H, and elements beyond the second period may have more than 8 electrons. It is important to remember that no elements of the second period may have more than 8 electrons! (Check back with number of bonds that elements can form). Why do we need a triple bond in N2? - to achieve octet.
Boron is an exception (incomplete octet): BH3 Phosphorus and sulfur are in the 3d period: PF5 (no NF5 but NF4+!) SF6 4. Lewis Structure Writing: 1. Count all the valence electrons available for all the atoms involved (for main group elements the # of valence electrons is the # of the group). If the particle you are trying to make has a charge, add or subtract the appropriate number. 2. Choose the central atom(s) and arrange the rest of them around it. Remember, hydrogen cannot be a central atom (it can only form one bond). In some cases it is difficult to decide which atom is central, and you should just know it, but in most cases you should be able to do this. 3. Make one bond between the central atom and each other atom. 4. Count the amount of electrons used at this point and left over. Remember, you can only use that total amount you have calculated in #1, neither more or less!
5. Distribute the rest of available electrons between atoms in the most symmetrical fashion and in the way that each atom (except for H) has 8. If there is not enough to satisfy all atoms, more electronegative atoms achieve octet first. 6. Remember that “octet rule” can be violated for the elements beyond the 2nd period (they have d-orbitals available to place extra electrons in). 7. Determine the Formal Charge on each atom: Formal Charge = # of valence e - (# of unshared e + 1/2 # of shared e) A better Lewis structure has no formal charges (or smaller charges). Examples: valence electrons #
CH4 NH4+ CO32H2SO4 HClO HNO3
4 + 4(1)= 8 5 + 4(1) – 1 = 8 4 + 3(6) + 2 = 24 2(1) + 6 + 4(6) = 32 1 + 7 + 6 = 14 1 + 5 + 3(6) = 24 Formal Charge: N: 5 – 1/2(8) = + 1
O H O N O
F F B F
Useful thing to remember about Formal Charges of “organic atoms”: - if carbon has 4 bonds, FC = 0 - if carbon...