Chapter 8 and 9 Study Guide

Chapter 8:  Ionic Bonding

Chapter 9:  Covalent Bonding

What is a Chemical Bond? The electrostatic forces holding atoms together to form a molecule (These Chemical bonds are called: Intramolecular Bonds).  These bonds can be a give-take relationship between two atoms (ionic bonds) or a shared electron relationship between two atoms (covalent bonds).

Ionic Chemical Bonds – If one atom desires to have an electron more than another, then when they get together,  they will form an Ionic Bond.  For example, Sodium will GIVE UP its valence electron to Chlorine so that Chlorine can have a full octet in its valance.

-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.

    Atoms that give up electrons become positive and are then called: “Cations”.

    Atoms that take electrons become negative and are then called: “Anions”.  

    Sodium becomes a Cation.  Chlorine becomes an Anion when they form NaCl.

-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 bond. 

-How do you know two atoms will form an ionic bond (as opposed to a covalent bond)?  Look at page 263 for the table of electronegativities.   Ionic bonds form when there is a greater than 1.7 electronegativity difference (p. 263-264)  between the two atoms.  Why?  Because that means that one atom REALLY wants the electron more than the other.  A difference of 1.7 or more makes the bond ionic (as opposed to covalent)

(Examples are:  NaCl, MgCl₂)

-The key to ionic bonding is getting two (or more) elements together to form a compound such that they all have an OCTET in their valence energy level.  This makes a stable structure and forms a chemical compound.

-This Sounds odd, but it’s True:  Cations are POSITIVE ions.  Anions are NEGATIVE ions.  However, if you have an electrode in a solution, the POSITIVE electrode is called an ANODE and the NEGATIVE electrode is called a CATHODE!  What the heck?  The reason why is because the POSITIVE ANODE attracts NEGATIVE ANIONS.  Anions go to the Anode.  Cations go to the Cathode.

Covalent Chemical 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. 263 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.

Non-polar covalent bonds  - zero – 0.3 electronegativity difference (p. 263-264) 

(Examples are: H₂, N₂, O₂)

Polar covalent bonds - between 0.3 and 1.7 electronegativity difference (p. 263-264) 

(Examples are: HCl, H₂O, CO₂)

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 9-17 p. 265)

A non-polar covalent compound would be like H₂, N₂, O₂…etc because they have no net electronegativity difference.

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) , Sn (6)

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. 256-257):  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 9-3 on page 260)  - I suggest strongly that you copy the chart onto a separate sheet of paper with the headings (below).  This way you can see how the “stick” figures (called Lewis structures) predict the VSEPR shape

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The Following Notes refer to Chapter 13.2 (page 393-395 of your textbook)

INTERMOLECULAR FORCE:  The force holding molecules together.  The phase of matter is based on this force – also an electrostatic force, like INTRAMOLECULAR, but much WEAKER! 

     Boiling point/Freezing point is based on INTERMOLECULAR FORCES.  For Example, you would not break the individual INTRAMOLECULAR 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 393-395 and below.

      London Dispersion forces p. 393:  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.  Non-Polar molecules have Dispersion forces only!

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 to Dispersion Forces:  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 (p. 394):  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 already 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.  Polar Molecules have Dipole-Dipole Forces

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

      Hydrogen Bonding forces (p. 395):  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.  Polar Molecules WITH HYDROGEN ATOMS have Dipole-Dipole Forces AND HYDROGEN BONDING!

     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.

     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!

In Summary:

   There are two kinds of Chemical Bonds:

1.  INTRAMOLECULAR  - These are the electrostatic bonds BETWEEN two atoms to make up a molecule.  They are VERY strong.  They come in either Ionic or Covalent form.

2.  INTERMOLECULAR – These are the electrostatic bonds BETWEEN two MOLECULES.  They are VERY weak!  They are responsible for holding together two or more molecules to make a compound (example, water is in liquid form because of INTERMOLECULAR bonds).

Polarity of Bonds:

      Whenever you have a bond, you have the possibility of Polarity (an un-equal sharing of electrons).

     There can be polarity between two atoms that make up a molecule

      (this is called Bond Polarity)

      There can be polarity of the entire MOLECULE

      (this is called molecular polarity).

You MUST at least have bond polarity in order to have an overall molecular polarity.

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

     Example:  A molecule with both 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, MOLECULAR polarity. (example:  Bent - H₂O; Trigonal pyramid – NH₃; and Trigonal planar – SO₂).

     Example:  A molecule with Bond polarity BUT 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).

     Example:  A molecule with 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)

In summary:

How do you know when a molecule has polar bonds?  Any two atoms bonded together, which are not the same atom, has polar bonds.  Example:  HCl has polar bonds.  H₂ does not.

How do you know when the molecule itself is polar?  1)  Any molecule with lone pairs of electrons on the CENTRAL atom is polar (if the lone pairs are only on the atoms around the central atom, then it isn’t polar).  2) Any molecule with different atoms around the central atom is polar (for example CH₄ has 4 hydrogens - same.  CH₃Cl has 3 hydrogens and 1 chlorine (chlorine makes it “different”).

Can a molecule have polar bonds but itself not be polar?  You bet!  Draw the structure of CO₂ or CH₄ and you’ll see that it fits the polar bonds model, but not the polar molecule model.

Can a molecule not have polar bonds but be polar?  No way!  You have to have some polarity within the molecular bonds in order for their to be an overall molecular polarity.