LEARNING OBJECTIVES:

• Ionic Bonds

• Recognizing ionic bonds
• Identifying ionic species

• Covalent Bonds
  
• Recognizing covalent bonds
• Identifying covalent species
  
• Writing Lewis Dot Structures
• Electronegativity and Polar Covalent Bonds
• Shapes of Molecules
  
• Molecular Structure
• VSEPR Model

 IONIC BOND

•  Ionic Bonds are strong, attractive, electrostatic forces between cations and anions

• substances which form ionic bonds, then, must consist of identifiable cations and anions
• the bonds are always formed between oppositely charged ions
  
• They are generally formed between metals (which typically form cations) and non-metals (which typically form anions)
• Overall, the formation of an ionic bond releases energy
• Ions form from atoms in order to achieve a noble gas configuration – an octet
  
• atom and ion symbols are shown as Lewis symbols (which show the valence electron configuration) such as for those atoms below:



• When a metal atom loses an electron to become a cation, a non-metal atom accepts the electron to become an anion

Examples of the formation of ionic compounds:

•  ionic compound NaCl is formed from one Na⁺ ion and one Cl^- ion:
  
  Na ® Na⁺ + 1e^-             this represents the formation of a Na⁺ octet
  
  Cl + 1e^- ® Cl^-                this represents the formation of a Cl^- octet
  
  Na⁺ + Cl^- ® NaCl
  

• Na loses one electron and Cl gains one electron as shown



• ionic compound MgCl₂ is formed from one Mg²⁺ ion and two Cl^- ions:
  
  Mg ® Mg²⁺ + 2 e^-           this represents the formation of a Mg²⁺ octet
  
  Cl + 1e^- ® Cl ^-                 times 2 represents the formation of 2Cl^- octets
  
  Mg²⁺ + 2Cl^- ® MgCl₂
  
• Mg loses two electrons, but each Cl can gain only one electron, so it takes two Cl to accept both Mg electrons



•  ionic compound AlCl₃ is formed from one Al³⁺ ion and three Cl^- ions:
  
  Al ® Al³⁺ + 3e^-         this represents the formation of a Al³⁺ octet
  
  Cl + 1e^- ® Cl^-           times 3 represents the formation of 3Cl^- octets
  
  Al³⁺ + 3Cl^- ® AlCl₃
  
• Al loses three electrons, but, again, Cl can gain only one electron, so it takes three Cl to accept all three Al electrons
  
• ionic compound Al₂O₃ is formed from two Al³⁺ ions and three O^2- ions:
  
  Al ® Al³⁺ + 3e^-           times 2 represents the formation of two Al³⁺ octets--loss of 6 electrons
  
  O₂ + 4e^- ® 2O^2-         times 11/2 formation of three O^2- octets--gain of 6 electrons
  
  2Al³⁺ + 3O^2- ® Al₂O₃
  
• Al loses three electrons, but O must gain two (to be a proper ion with a full octet), so it takes two Al to lose six electrons and three O to accept the six electrons
  
• Notice that in the formation of an ionic compound, the number of electrons lost in the formation of the cation(s) must equal the electrons gained in the formation of the anion(s). We cannot have "loose" electrons floating around!
  

 COVALENT BOND

•  A covalent bond is a pair of electrons shared between two atoms

• Generally occurs between two non-metals in order that each atom can attain an octet of electrons through sharing
• the bonds are always formed between like atoms that are neutral (not charged)
• They are generally formed between non-metals

• Covalent bonds typically release energy upon formation
• Two atoms can share one pair, two pair or three pair of electrons to form a single, double or triple bond, respectively
• Bonding Pairs refer to any electron pairs shared by two atoms in the formation of a covalent bond
• Nonbonding Pairs or Lone Pairs refer to any electron pairs which are not shared by two atoms and reside only on one atom

 examples:

 H₂

 

Notice that H₂ has one pair of shared electrons (a single bond) between the two H atoms and it has no unshared (nonbonding or lone pairs) electrons.

O₂

 

Notice that O₂ has two pair of shared electrons between the two O atoms or a double bond, and two pair or four nonbonding or lone pair electrons on each O atom

N₂

Notice that N₂ has three pair of shared electrons between the N atoms or a triple bond and one pair of nonbonding electrons on each N atom.

 NH₃

Bonding Pairs: any electron pairs shared by two atoms in the formation of a covalent bond
Non-bonding Pairs or Lone Pairs: any electron pairs which are not shared by two atoms and reside only on one atom



LEWIS DOT STRUCTURES 

• Used for the prediction of bonding structures in molecules
• Used to show the relative position of all electrons – bonding and nonbonding – in a molecule
• All atoms should have an octet of electrons (except H which will only have 2)
• All electrons should be paired
• Rules

1. write Lewis symbols for all atoms in the molecule
2. count the total number of electrons
3. count the number of unpaired (single) electrons
4. divide the number of unpaired electrons by 2 – this is the number of covalent bonds in the molecule
5. arrange the atoms appropriately: more metallic element is in the center, the less metallic elements are pendant atoms (hydrogen is never a central atom)
6. put the covalent bonds in place
7. place all remaining electrons around atoms such that all atoms have an octet
8. count all electrons again – should be the same as #2 – make sure all atoms have an octet (all shared electrons are counted in the octet for both atoms sharing the electrons

examples:

• example with single bonds:



• notice that there are three bonding pairs and one non-bonding pair of electrons and the hydrogens are terminal atoms (never in the middle of a Lewis Dot Structure)
• also notice that we typically represent a single bonding pair of electrons with a straight line between the two atoms 

• example with double bonds:

• Since there are only two places to put bonds, all four covalent bonds must occur between the three atoms--either as two double bonds, or one single bond and one triple bond
• Notice that the central atom is the more electropositive atom (the atom more to the left on the periodic table) while the "pendant" atoms are the more electronegative atoms (atoms more to the right on the periodic table)
  
  

 

• Example with a combination of single and double bonds:
  
  

• Example of a triple bond

Structures of Cations and Anions

When writing Lewis Dot Structures of anions or cations, the rules are the same except for the addition of rule #2a: if the ion is an anion, add the number of electrons corresponding to the charge of the anion, if the ion is a cation, subtract the number of electrons corresponding to the charge of the cation. All the other rules still apply.

• Example of an anion with single and double bonds:

BOND POLARITY

• electrons shared between two atoms with different EN are shared unequally
• unequal sharing creates a separation of charge -- polarity or a polar bond
• the greater the difference of the electronegativity of the two atoms in a bond, the greater the polarity (charge separation) of the bond because the electrons are more attracted to one atom that the other
• electrons shared between two atoms with the same electronegativity are shared equally and there is no charge separation, no polarity
• if there is no difference in the elctronegativity of the two atoms in a bond, then the bond is a nonpolar bond

Electronegativity

• A measure of the ability of an atom in a covalent bond to draw bonding electrons to itself
• Determines the direction of polarity in a covalent bond
• Electronegativity increases from left to right and bottom to top on the periodic table – the most electronegative is at the top right and the least electronegative is at the bottom left

examples:

Remember that electrons are not the little dots, as we have depicted them in the preceding diagrams, but rather "clouds" of negative charge. The electron "cloud" of electron density "hovers" more closely to the more electronegative atom, giving that atom a more negative charge (a partial negative charge, d-) and leaving the other atom exposed with a more positive charge (a partial positive charge, d+) as depicted below:

Problem:

• Which is more polar, HF or HI?
  the EN of H = 2.1 and the EN of F = 4.0 and the EN of I = 2.5
  the EN difference for HF is 1.9 and the EN difference for HI is 0.4
  the greater the EN difference, the more polar the bond, so HF has the more polar bond
  
• Describe the "ultimate" polar bond.
  The "ultimate" polar bond is the one with the greatest charge separation--the ultimate in charge separation occurs when one atom completely loses control over its electron(s) and the other atom completely takes control of the electron(s). This would, in effect, be an ionic bond.

SHAPES OF MOLECULES

• The reactivity of a molecule is determined, in large part, by the shape of the molecule
• The shapes of molecules are determined by the relative position of the atoms
• In turn, the position of the atoms is determined by the number and positions of the electron pairs (both bonding and nonbonding) that is why it is so very important to know how to predict Lewis Dot Structures of molecules
• the normal distance between the atoms in a covalent bond is called the bond length and it essentially measures the "overlap" of the electron clouds of the two atoms participating in the covalent bond

VSEPR

Valence Shell Electron Pair Repulsion model is a model used to predict the relative position of electron pairs residing around a central atom. It is based on the premise that all electron pairs residing around an atom wish to be as far away from one another as possible and as symmetrically located relative to one another as possible. For example:

• 2 electron pairs residing on an atom will arrange themselves 180° apart – molecule is linear
• 3 electron pairs – 120° apart – molecule is trigonal planar
• 4 electron pairs -- 109° apart – molecule is tetrahedral

If there are nonbonding electrons around the central atom, the shape that the atoms of the molecule describe will be different than the shape that the electron pairs of the central atom describe

There is an excellent site on the world wide web that has a tutorial on VSEPR as well as writing Lewis Dot Structures. In order to see and manipulate the 3D models presented, you need to download the Chime program (for more information on that subject, go to the tutorial and click on Mark Winter's VSEPR page link at the bottom). I do not require that you use the tutorial, but I strongly encourage you to check out the site of the VSEPR Tutorial . There is also a link from my homepage.

• Figure I: Central atom with two bonding electron pairs:

Figure II: Central atom with three bonding electron pairs:

• Figure III: Central atom with four bonding electron pairs:

In the three diagrams above, the electron pairs describe the relative geometry, but we've not indicated which pairs might be bonding pairs and which might be nonbonding pairs.

Look at Figure II. If all three pairs of electrons were in fact bonding pairs, the shape of the molecule (as indicated by the relative position of the atoms attached to the central atom) would be trigonal planar--the same as the geometry of the electron pairs. But what if one pair of electrons is a nonbonding pair (remember, we can't "see" a pair of electrons--we look for covalently bonded atoms to let us know that the electron pairs are there). What would the shape (as defined by the relative positions of the atoms) of the molecule be? It would not be trigonal planar, because there are only three atoms, the central atom and two pendant atoms--but we still have three electron pairs which arrange themselves in trigonal planar geometry. The shape that the atoms present, however, would be bent.

• Figure IV: Central atom with no lone pairs and with one lone pair:

• Figure V: Central atom with no lone pairs, one lone pair and with two lone pairs:

Polar Molecules

When molecules contain polar bonds they may or may not be polar molecules – it depends upon the relative positions and symmetry of the polar bonds

examples: