Chapter 6 Study Guide - Chemical Bonding


Chemical Bond – forces holding atoms together to form a molecule (Intramolecular)  By examining the forces (electronegativity difference – see p. 151) between atoms in a compound, you can determine if a molecule is covalent or ionic:  The three types of bonds are:

Non-polar covalent bonds  - zero – 0.3 electronegativity difference (p. 163)  (H2, N2, O2)

Polar covalent bonds - between 0.3 and 1.7 electronegativity difference (p. 163) 

(HCl, H2O, CO2)

Ionic bonds - greater than 1.7 electronegativity difference (p. 162)  (NaCl, MgCl2)


Ionic Bonds - Cations are Positive ions, Anions are Negative ions – Wow, great general question!

                 -Cathodes are negative electrodes, Anodes are positive electrodes – great general question!

             -Ionic compounds exist as crystal lattice structures.  One atom donates its

electron to another atom and they “hang out together” because of the positive/negative attraction.

-Metals and non-metals often react to form ionic compounds.  If the

two elements are from opposite sides of the periodic table, it is most likely an ionic

             compound (NaCl is a good example)

Rules/Tips for Drawing Lewis Structures:  (IONIC MOLECULES ONLY) 1.7 or greater

1.  The first element is positive and the second one is negative.  Look at the column the first element is in and put a plus charge next to it as needed.  (see examples to the right)

2.  The second element is negative and it has a full octet of electrons surrounding it.  Put it in brackets and put a negative charge next to it as needed. (see examples to the right)

Covalent Bonds - A sharing of electrons is called a covalent bond.  If the electronegaitvity difference between two atoms is 1.7 or less, then it is probably a covalent bond.  Look on p. 151 for a table of electronegativities.  There are two types of covalent bonds:  Non-polar covalent and polar covalent.  Use the electronegativity chart to determine which one you have.

Polar vs. Non Polar covalent bonds:  A polar bond is a bond in which the electron(s) tend to favor one atom over another.  For example, in HCl, the electron would prefer to be with Cl than with H.  Therefore, the Cl is slightly more negative and the H is slightly more positive.

This is called a dipole moment and is indicated by a partial positive or a partial negative sign.  For example, you could indicate the dipole moment on HCl as follows:  HCl (fig 6-3 p. 163)

A non-polar covalent compound would be like H2, N2, O2…etc because they have no net electronegativity difference. (fig 6-3 p. 163)

Rules/Tips for Drawing Lewis Structures: (COVALENT MOLECULES ONLY) 1.7 or less

1.  Count up the TOTAL valence electrons for your molecule or polyatomic ion (remember to ADD electrons if it has a negative sign and to SUBTRACT electrons if it has a positive sign).

2.  Draw a single bond between the central atom and all of the atoms to which it is attached.  The central atom is USUALLY the first atom in a compound, but it is NEVER Hydrogen.  (Remember that each single bond counts as TWO electrons.)

3.  After you have drawn your single bonds, put the rest of the electrons (from the total number of valence electrons) around your atoms in your molecule.  Remember that each atom must have 8 electrons around it or shared between it and another atom.  (There are, however, exceptions to the octet rule.  H, for instance, needs only 2 electrons.  See below for other exceptions)

Exceptions to Octet Rule:  H (2e),  Be (4e),  Al  (6e),  B (6e) ,  Ga (6e)

4.  Remember that if you see a double bond, that means that 4 electrons are being shared between those two atoms.  If you see a triple bond, that means there is 6 electrons being shared.

5.  No atom may have 1 electron on any one of its sides.  Electrons travel in pairs, you will never find an orbital with only one electron in a molecule, it must always have a completely filled or an empty orbital.


Resonance (p. 175):  If a covalent molecule requires a double bond on one side, but not on the other of a central atom, that double bond may “resonate” throughout the molecule.  Take the example of SO2.    


Molecular shapes  (see table 6-5 on page 186)  - Memorize this chart

Shape      What it looks like      Atoms bonded to central atom       Lone Pairs          Bond Angle

















INTRAMOLECULAR FORCE:  The force holding individual atoms together to make

             up one molecule.  Molecular shapes are based on this force.  Whether a molecule is ionic

             or covalent.  Whether a molecule is polar or non-polar. 


Predicting the polarity or non-polarity of a molecule by looking at structure

Bond polarity and Molecular polarity:  Any molecule which has a central atom and some unbonded electron pairs around it.  This provides UN-equal pulling of electrons and thus, polarity. (example:  Bent - H2O; Trigonal pyramid – NH3; and Trigonal planar – SO2).

Bond polarity and No Molecular polarity:  Any molecule which has a central

atom and NO unbonded electron pairs around it.  This provides equal pulling of

electrons and thus, no polarity. (example:  Linear – CO2;  Trigonal planar – BF3; Tetrahedral – CH4).

            No Bond polarity and No Molecular polarity:  Any Diatomic (N2, H2, O2…)

A Molecule cannot have molecular polarity unless it also has bond polarity.  But a molecule with bond polarity does not necessarily have to have molecular polarity. (see p. 190-191 if you don’t understand this comment)


INTERMOLECULAR FORCE:  The force holding molecules together.  The phase of

matter is based on this force.  Boiling point/Freezing point.  When you boil or freeze a substance, the boiling point and freezing point are dependent upon the substance’s INTERMOLECUAR forces only.  For Example, you would not break the individual bonds between hydrogen and oxygen when you melt ice.  You would merely break the bonds between water molecules.


There are three basic INTERMOLECULAR FORCES (also called Van der Walls Forces).  They are described on pages 189-193 and below.

                        London Dispersion forces p. 193:  depending upon where an electron is in its orbit around the nucleus, a “pole” may be set up.  Wherever the electron is would be more negative than where it isn’t.  Example:  Hydrogen is a diatomic atom and it has 2 electrons.  Hydrogen also has an electronegativity difference of 0 (non-polar covalent).  If the electrons go to one side of the molecule, then the molecule has a positive and a negative end which would then be attracted to other positive and negative ends of other hydrogen molecules floating by.

     Dispersion forces occur between two otherwise nonpolar molecules.  Dispersion forces are the weakest of all three van der Waals forces.  All molecules, including polar ones, have dispersion forces. As a rule: As molar mass increases, so do the dispersion forces (because there are more electrons) and so do the boiling points.

     Bottom line:  Large dispersion forces make molecules want to “hang-on” to each other so it takes more energy to boil them away from each other.  Similarly, low dispersion forces would make molecules not want to get together so it would take a good deal of COLD to get fluorine to freeze as opposed to say Iodine.  Dispersion forces are the only way non-polar molecules stick.

     Noble gases only have dispersion forces and this is why they must get so COLD to condense or freeze.

                        Dipole-Dipole forces:  Like dispersion forces, these are created by electrons being on one end of a molecule.  Unlike dispersion forces, dipole-dipole forces are on polar molecules.  This means that these molecules all ready have those electrons shoved to one end of the molecule.  An example is PCl3.  This Trigonal Pyramidal molecule has a lone pair of electrons on the P atom and is therefore a polar molecule.

     Bottom line:  The more polar you are, the higher your freezing point and boiling point.  Molecules with dipole-dipole forces also have dispersion forces.


                        Hydrogen Bonding:  Hydrogen has an electronegativty number of 2.1.  When it is attached to Nitrogen (3.0), Oxygen (3.5) or Fluorine (4.0), the electron is drawn more to those atoms and thus Hydrogen is left with a partial positive charge and they with a partial negative charge.  A good example is our old polar friend dihydrogen monoxide:  H2O.

     Since hydrogen is so small also, it tends to hug up next to the negative charged N, O, or F of another molecule.  This is called hydrogen bonding (see page 192 fig 6-28.)

     Carbon has a electronegativity of 2.5 and is therefore close to that of hydrogen so the hydrogen bonds are not as strong.

     Hydrogen bonding is about 10 times a strong as regular dipole-dipole forces.  Therefore, a molecule of HF has an incredibly high boiling point, and so does water.  Molecules with hydrogen bonding also have dipole-dipole and dispersion forces.

      NOTE:  H2 does not exhibit hydrogen bonding.  Only Dispersion forces.  Why?  Because Hydrogen-Hydrogen bonds have equal electronegativity pulls so there is no dipole effect on diatomic hydrogen.  WATCH OUT, THIS IS OFTEN A TRICK QUESTION ON TESTS!  Just because it is hydrogen doesn’t mean it will form an intermolecular hydrogen bond!


Homework for this chapter is worth 15 points

Instructions:  I will assign each homework as we go along.  Make sure that you indicate which homework part you are doing on your paper (sub-headings!!!!!)

      I will collect homework 1-4 on the day of the quiz.  Be sure to do ALL of the problems – both the ones from the book and the ALSO ones I wrote.


Homework Part 1: Is it ionic or covalent bonding?  Electron dots for elements

Section review p. 163 #3; Problems #33 & #37 p. 196


Homework Part 2:  Electron dot diagrams for covalents and ionics

Practice #1-2 p. 172, Practice #1-2 p. 174, Section Review #4 p. 175 (these are all covalent)  Problem #38 p. 197 (ionic), Problems #39, 41 & 42 p. 197 (all covalent)

Also:  Draw the Lewis dot structure for the following ionic compounds:

Aluminum Chloride

Potassium Chloride

Sodium Oxide

Aluminum Sulfide


Homework Part 3:  Predicting the Shape of a molecule using VSEPR

Practice p. 185 #1, Practice p. 187 #1, p. 196 #24, P. 197-198 #43, 45, 46, 48, 49, 51, 57

ALSO  draw the structure and name the shapes for the following molecules:

            H2, H2O, HBr, CH3NH2 (Where C is bonded to N and 3 hydrogens.  N is bonded to C and 2 hydrogens), CO2, H2CO, C2H2, C6H6, CH3Cl, NH3, CH4, O2, AlH3, CH3OH


Homework Part 4:  Intermolecular forces and determining polar/non-polar molecules.

p. 197 # 47.  Also do problem #48 only tell me if the bonds of those are polar or non-polar.

ALSO Look at the following molecules:  , CH3Cl,  O2  NH3  H2O    PCl3     BF3  H2

-Which ones have polar bonds?

-Which ones have molecular polarity?

-Which ones have ONLY dispersion forces?

-Which ones have dipole-dipole forces?

-Which ones have hydrogen bonding forces?