The more correct way to do Lewis Counting Valence Electrons Dot Structures (book method) Lone pair electrons belong entirely to  1. Get the sum of all valence electrons from the atom in question. all atoms. Ignore which electrons came Resonance Structures  from which atom. Shared electrons are divided equally  2. Arrange the elements between the two sharing atoms.  3. Place the electrons anywhere in the  (Valence electrons)assigned = compound to satisfy the octet and duet rule.  ( # lone pair electrons) +  1/2 (# of shared electrons).

Rules Formal Charge Formal charge To determine the formal charge on each The sum of the formal charges of all atom, take the number of valence atoms in a given molecule or ion must  Most molecules have a formal charge of 0 on all atoms. electrons assigned to the atom in the equal the overall charge on that species. molecule and subtract if from the number  If different Lewis structures exist of valence electrons on the free, neutral for a species, those with formal charges atom. on all atoms closest to zero are the best. Formal charge = valence electrons of the All negative formal charges should be  H 1 electron on free – 1 assigned = 0 free atom- valence electrons assigned on the most electronegative atoms.  O 6 electrons on free – 6 assigned = 0

HClO Ions The one on the right is correct

 For ions, an electron is added for each negative charge and an electron is subtracted for each positive charge from the total valence electrons.  All ions will have a formal charge on at least one atom

 Na+ would subtract 1 electron for the +1 2-  SO4 would add two electrons for the 2-

1 Molecules/ions with a formal Resonance Lets look at ozone charge

 O3  A very important point of this method is showing that electrons don’t “belong” to It may look like this Or it may look like this any atom in a molecule.  CO  The electrons can “flip” places.

-  NO3 Data suggests it looks like both at the same time. These are called resonance structures,  CO 2- 3 which are all possible Lewis Dot structures for a molecule.

What this means Stability

 Double bonds and a single bond are different  The resonance structure makes it like there lengths. are 2 “one and a half bonds” instead of 1 Differences between Covalent  Looking at ozone, you would expect one single and 1 double bond. Bonding and Ionic Bonding oxygen to be closer to the middle oxygen than  This makes compounds much more stable the other. or non reactive.  Experiments put their bonds at the same length that is somewhere in between the length of a single and double bond.

Shortcut to determining type of Major difference bond Why this works  Covalent bonding is a sharing of electrons, Ionic  When a metal and nonmetal bond you get an bonding is a transfer of electrons. ionic bond  Electronegativity- ability of an atom to attract  Covalent bonds are between a small number of  ~ something from the left excluding H bonds and hold bonding electrons. atoms. with something from the right = ionic bond.  Elements with a large difference in  H-O-H one oxygen bonding to two hydrogens  When two nonmetals bond you get a covalent electronegativity will form an ionic bond,  Ionic bonds are between a very large number of bond elements with a small difference will form ions stuck together.  ~things from the right bond with each other covalent bonds.  In NaCl every sodium ion is bonded to every =covalent bond. single chloride anion near.  Metals don’t bond with each other.

2 Using the periodic table to Electronegativity Chart What about the middle ground? determine electronegativity  electronegativity generally increases  What if the difference in electronegativity up and to the right excluding noble isn’t large or small but in the middle? gases.  For example H (2.1) and O (3.5)  Fluorine is the most electronegative element  These elements form a polar covalent bond. (4.0) followed by oxygen (3.5) and chlorine  Polar Covalent Bond- unequal sharing of (3.0). electrons in the bond  A full chart is on page 344.  so the electrons stay around oxygen more than hydrogen

Do any bonds have an equal Polar covalent Bonds sharing?  4 electrons occupy  Yes, (normally the same element) when

this cloud. elements are equally electronegative like O2  Notice how much  In fact, anything with a very slim difference larger the cloud is (less than 0.5) in electronegativity will around oxygen as pretty much equally share electrons. compared to  Nonpolar covalent bonding- equal sharing hydrogen. of electrons in a bond

Why it is called polar Denoting positive and negative Dipole Moment  polar implies different ends have different charges Neither side is completely positive or negative, Dipole moment- property of a molecule similar to a magnet. they are only partially positive and partially negative. where the charge distribution can be HCl as a polar bond, meaning represented by a center of a positive charge the electrons stay around Cl more than H and a center of negative charge.  It is represented by this symbol

That makes this and this side Chlorine H + side negative positive - Chlorine H  Positive center Negative center

3 So the dipole moment for HCl… is represented like this. Note the positive charge is on the hydrogen side, the negative on the chlorine side.

H

Chlorine

4