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Need STRUCTURES for Covalent Molecules

Need STRUCTURES for Covalent Molecules. Because there are no ions in covalent molecules, we can’t use charge to predict the ratio of elements in a molecule. Actually, most nonmetals can bond covalently in many possible ways.

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Need STRUCTURES for Covalent Molecules

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  1. Need STRUCTURES for Covalent Molecules • Because there are no ions in covalent molecules, we can’t use charge to predict the ratio of elements in a molecule. • Actually, most nonmetals can bond covalently in many possible ways. • Instead, you will be told the ratio for a particular molecule by a chemical formula or name. • However, it turns out that the specific arrangement of atoms in a molecule is VERY important to its behavior: • It determines the molecule shape • It determine the distribution of electrons (polarity) • Which in turn determines how it will interact with neighbor molecules

  2. Electron-Dot Symbols are a symbolic way to show an atom’s valence electrons to help us see how it might bond with a neighbor atom • Represent the valence electrons as dots around the four sides of the chemical symbol for the atom • Each side should have one dot before any side gets two dots • can be drawn in more than one way Magnesium has 2 valence electrons (12 e-: 2 in 1st level, 8 in 2nd level, 2 in 3rd level)

  3. Electron-Dot Symbols for Selected Elements

  4. How to draw a Lewis Structure – method 1(for a covalent molecule)

  5. Examples – CO2, CClN C N Cl C O O C N Cl C O O C N Cl C O O

  6. Terminology this is a double bond (2 pairs of valence electrons – two from each neighbor atom - shared between atoms) this is a single bond (a pair of valence electrons – one from each neighbor atom - shared between atoms) C N Cl C O O this is a triple bond (3 pair of valence electrons – three from each neighbor atom - shared between atoms) this is a lone pair (a pair of valence electrons not shared between atoms)

  7. How to check a Lewis Structure • Electron count: the total count of electrons must equal the sum of all the valence electrons for the atoms in the molecule • Octet check: all non-H atoms have an octet (8 electrons in the form of bonds (2 each) and lone pair (2 each). H atoms should have a duet (2 electrons in the form of 1 bond) • 4 bonds or • 3 bonds + 1 lone pair, or • 2 bonds + 2 lone pair, or • 1 bond + 3 lone pair. • Valence / formal charge check: The atoms' original valence should be preserved (unless the electrons were pushed*). To count valence after its bonded, we only count half of the bonded electrons (so 1 for every bond), but 2 for each lone pair. An easier way to check is that • Carbon (all Group 14) should have 4 bonds and no lone pair (= 4 valence e) • Nitrogen (all Group 15) should have 3 bonds and 1 lone pair * (= 5 valence e) • Oxygen (all Group 16) should have 2 bonds and 2 lone pair * (= 7 valence e) • Fluorine (all Group 17) should have 1 bond and 3 lone pair (= 7 valence e) • Hydrogen should have 1 bond and no lone pair (= 1 valence e) *Group 15 & 16 occasionally don’t have the proper valence because electrons were pushed around. They still need to have an octet, though. And the total electron count must be correct.

  8. Example: When an H atom is near another H atom, they bond in pairs: H H H H H H H H H H H H H H H H H H As found in nature Similarly: F F Br Br O O N N Cl Cl I I

  9. The “Diatomic Seven” • Seven elements (all with high electronegativities) are so well suited to bonding with themselves that they are never found as atoms, but always as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2 • (the non-H elements sort of form the number 7 on the periodic table to help you remember them) • When you see the chemical names “hydrogen”, “nitrogen”, “oxygen”, “fluorine”, “chlorine”, “bromine” and “iodine” you need to know that their chemical formulas are H2, N2, O2, F2, Cl2, Br2, I2 • (to refer to individual atoms of these elements that are not found in nature, we say “atomic hydrogen”)

  10. Front table

  11. # 3 is ADVANCED! First form single bonds between the N and one O Then form double bonds between the N and another O Final picture should show charge: 1- O O O O O N N N N O O O O O O O O O O Now add the extra electron to the O with a single bond N Now PUSH one unpaired O electron to the other to make a lone pair, And PUSH the lone pair of electrons on N to be between the N and O and make the bond

  12. # 3 another method Total up all electrons (5+6+6+6+1 from charge) = 24 Draw N in middle and O’s around (no e-’s) Connect with 1 bond each to center, which takes up 6 electrons O O O O O O O O O (There are 18 e- left to distribute) Because O is more electronegative, we will fill its octets until we run out N’s octet is NOT complete, so PUSH one lone pair of an O to make a double bond with N 1- O O N N N N Final picture should show charge: O

  13. # 3 brings up another phenomenon called resonance: We randomly picked the O that made the double bond. It could have been any of the 3 O’s. Really, I have three equally valid versions of this structure. Each of these equally valid pictures is called a resonance structure (we show double arrows between them to indicate the resonance) O O O O O O O O O 1- 1- 1- N N N

  14. Predicting Shape (of a covalent molecule)Valence Shell Electron Pair Repulsion theory Electron domains are areas of electrons that must remain together. A bond (single, double or triple) is 1 electron domain around a central atom A lone pair of electrons is 1 electron domain around a central atom This chart will help us to predict the shape of a molecule from its Lewis Structure

  15. Polar Bonds • In a covalent bond, even though the two atoms are SHARING the pair of electrons, if they don’t have the same electronegativity (pull on e–): • The more electronegative atom will pull the pair of electrons more and thus have a partial – charge (-) • The less electronegative atom will have a partial + charge (+) • If the DIFFERENCE in electronegativity between two atoms is > 0.45 but < 1.7, Dr. Little considers the bond to be polar and covalent • If the DIFFERENCE in electronegativities is > 1.7 there is enough difference for the more electronegative one to completely pull the electron away from the other atom, and thus it is considered an IONIC BOND. • It may be helpful to consider a tug-of-war. The flag (- charge) will be closer to the team that can pull harder.

  16. Drawing a polar bond Not necessary, but sometimes it helps visual people: • Determine dipole or polarity by subtracting electronegativities (3.0-2.5 = 0.5) • Draw an arrow from lower EN atom to higher EN atom and make a cross on the non-arrow end • You can also add the actual EN difference on top of the arrow • Instead of the arrow, you can also use + on the lower EN atom and – on the higher EN atom to indicate the partial negative charge on the more EN atom and partial positive charge on the less EN atom 0.5 + -

  17. Dipoles (or Polar Molecules) • If the electrons in a molecule are not distributed evenly, and one side of the molecule is slightly – (-) while the other side is slightly + (+), we call the molecule polar or a dipole • To be a polar molecule, TWO conditions must be met: • The molecule must have at least 1 polar bond • The geometry of the molecule must not have cancelling polar bonds • A linear molecule with 2 identical polar bonds • A trigonal planar molecule with 3 identical polar bonds • A tetrahedral molecule with 4 identical polar bonds • Polar molecules do not have full charges, but partial charges. The partial charges are capable of attracting charges in other nearby molecules, causing them to stick together somewhat. (not as strong as ions clustering together, but strong enough to affect behavior)

  18. Examples C-O bond is polar (3.5-2.5=1.0) Shape is linear(central C has 2 bond domains + 0 lone pair domains) The molecule CO2 is nonpolar by symmetry (linear with 2 identical bonds) I-Cl bond is polar (3.0-2.5=0.5) Shape is linear (only 2 atoms) The molecule ICl is polar because of the polar bonds. 0.5 1.0 1.0 0.5 C-Cl bond is polar (3.0-2.5=.5) Shape is tetrahedral (central C has 4 bond domains + 0 lone pair domains) The molecule CCl4 is nonpolar by symmetry (tetrahedral with 4 identical bonds) H-O bond is polar (3.5-2.1=1.4) Shape is bent (central O has 2 bond domains + 2 lone pair domains) The molecule H2O is polar because of the polar bonds. 0.5 0.5 1.4 1.4 0.5

  19. Flowchart for deciding if a molecule is polar or not Draw a Lewis structure Lookup the electro-negativity for each of the 2 atoms in the bond; subtract them; if the difference is: 0-0.44 nonpolar 0.45-1.7 polar > 1.7 ionic For each bond in the structure, determine if the bond is polar Does the molecule have 1 or more polar bond? NO YES Count the # of bonding domains and the # of lone pair around the CENTRAL ATOM and lookup shape on VSEPR chart nonpolar Determine the shape Is it tetrahedral? NO YES Is it trigonal planar? Does the central atom have 4 identical bonds? NO YES NO YES Does the central atom have 3 identical bonds? Is it linear? polar nonpolar NO YES NO YES polar Does the central atom have 2 identical bonds? polar nonpolar NO YES polar nonpolar

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