Properties depend on the nature of the bonding between elements in the compound and the strength of these bonds. In 1916 G.N. Lewis suggested that covalent compounds are the result of electrons that become paired and share the space between the bonded atoms in molecules. He developed a simple representation of bonding in molecules. This representation is called the LEWIS ELECTRON-DOT MODEL Chemical Bonding: Classical Representation

The Lewis model begins by recognizing that not all electrons in an atom participate in chemical bonding. Electrons occupy a set of SHELLS or ORBITALS surrounding the nucleus. Electrons in INNER shells, called CORE electrons, do not significantly participate in bonding. The outermost, partially filled shell (called the VALENCE shell) contains VALENCE electrons that are directly involved in bonding.

With the exception of Helium, for the Group IA to VIIIA the number of valence electrons in a neutral atom is equal to the element’s group number.

To determine the number of valence electrons and core electrons Total number of electrons = atomic number Number of Valence electrons for the A block equals the group number Core electrons = Atomic Number – # of Valence electrons For example: Cl atomic number = 17 Total number of electrons = 17 Number of valence electrons = 7 (since Gr VII A) Core electrons = = 10 Lewis diagrams involve just the valence electrons and assumes that the core electrons are not involved in bonding and reactivity.

LEWIS DIAGRAMS Lewis diagrams provide useful predictions of bonding in molecules, and hence the molecule’s structure and reactivity

The Lewis dot diagrams for the first 10 elements are: To construct a Lewis diagram dots are used to represent each electron in the valence shell of the atom. HHe LiBeBCNOFNe For elements with less than four valence electrons the dots (electrons) are displayed singly around the four sides of the atoms (with the exception of Helium). For elements with greater than 4 valence electrons, pair electrons

The maximum number of valence electrons for the “A” block is 8 Atoms with this full set of valence electrons are particularly stable and hence unreactive. These elements are called the NOBLE GASES - He, Ne, Ar, Kr, Xe, Rn The noble gas elements exist naturally as monoatomic species and compounds containing these atoms are not particularly stable.

OCTET RULE - atoms tend to gain, loose, or share electrons until they are surrounded by EIGHT valence electrons. Hence, atoms give up, take up, or share electrons to gain an OCTET of electrons, and hence a stable electronic configuration (the noble gas electronic configuration). The octet rule applies to the A block elements. For the transition elements, including the lanthanides and actinides, the octet rule generally does not apply.

IONIC COMPOUNDS metallic element + nonmetallic element -> ionic compound Formation of NaCl Na --> Na + + 1e - Cl + 1e - --> Cl - Na + + Cl - --> NaCl Atomic number of Na = 11 Number of electrons in Na = 11 Total number of electrons in Na + is 10 The electronic configuration of Na + is the same as Ne

Cl has seven valence electrons. If it accepts 1 electron from Na, then Cl - has eight electrons completing an octet (same electronic configuration as Ar) Lewis diagrams for formation of NaCl ionic bond Na--> + 1e - Na + Cl+ 1 e - -->Cl - Na-->Cl+ Na is electropositive, Cl electronegative; Na transfers an electron to Cl Cl - Na +

LEWIS DIAGRAMS FOR COVALENT COMPOUNDS Nonmetallic element + nonmetallic element -> covalent compound Formation of Cl 2 Cl + Cl  Cl 2 -->Cl+ Lewis diagram By forming the Cl 2 molecule, each Cl atom has an octet of electrons, with two electrons shared by both Cl atoms

Lewis structures for covalent compounds show the shared pairs and the unshared pairs of electrons for each atom. The electrons shared in the bond are called the BONDING electron and those unshared the NON-BONDING electrons, or LONE PAIRS Cl bonding electrons non-bonding electrons

The bonding pair is usually represented by a line joining the two atoms. A single line represents a SINGLE bond between the two atoms (multiple bonds between atoms do form) NHH H ammonia Cl Hence for Cl 2, the Lewis dot structure is written as: NHH H

Covalent bonds can also be formed by the sharing of an electron pair supplied by only one of the two atoms. Example NH 3 + H + N H H H + H + --> NHH H H + ammonium ion The N atom in NH 3 possess a nonbonding electron pair that it shares with the H + ion, which has no electrons. Once this bond is formed, there are four N-H bonds; all four are indistinguishable.

Molecules with Double and Triple Bonds Bonds involving more than one shared pair of electrons are called MULTIPLE bonds. H2CH2CCH 2 H2CH2C Ethylene HCCH HCCHAcetylene NNNitrogen NN Examples

Guidelines for drawing Lewis Structures Lewis structures, while simplistic representations of molecules, do help predict chemical properties of molecules For example, the position of multiple bonds in a molecule affects it’s geometry and reactivity. Before drawing the Lewis structure, information on the molecular skeleton of the molecule must be known, i.e. which atoms is bonded to which.

Knowing the molecular skeleton, the guidelines below help draw Lewis structures for molecules. 1) Each element tends to achieve a noble gas electron configuration; this means two electrons for H and eight valence electrons for the other elements 2) H is always bonded with one electron pair (a single bond); H atoms are always at the end or terminal position in a molecule (molecules with the skeleton C-H-C, for example, do not exist). 3) O in compounds typically involves two shared electron pairs, C almost always four shared pairs and N normally involves three shared pairs. 3) Only a few elements commonly form multiple bonds (C,N,O,P) of which C and N can form double and triple bonds 4) Many molecules have a central atom about which other atoms are symmetrically arranged (e.g. CH 4 )

Steps for drawing Lewis diagrams - e.g. for azide ion, N 3 - 1) Count the TOTAL number of VALENCE electrons available by adding the group numbers of ALL elements in the molecule. If the molecule is positively charged, subtract the charge from the total number of valence electrons and if negatively charged add the charge. For N 3 - : Number of electrons available for bonding = 15+1=16 Call this number A; hence A = 16 2) Calculate the total number of electrons needed for each atom to have its octet satisfied. Call this number N Each of the nitrogen atoms need 8 electrons Hence N = 3 x 8 = 24

3) Subtract A from N. This is the number of SHARED electrons; i.e. bonding electrons, S S = = 8 4) Assign two bonding electrons (as one shared pair) to each bond between two atoms N-N-N The single bonds between the three N atoms account for 4 electrons. 5) If after assigning single bonds between atoms there are electrons left over assign multiple bonds. In some cases there maybe more ways to do this, but in general, double bonds form between atoms of C,N,O & S. For N 3 - assign double bonds between the three N atoms N=N=N

6) Assign the remaining electrons as lone pairs to the atoms so that each element satisfies the octet rule (two electrons for H) N=N=N The Lewis diagram for N 3 - is: Count the electrons for each atom to make sure each satisfies the octet rule. Count the total number of electrons to make sure the Lewis diagram agrees with the total number of electrons available For the diagram above, number of electrons = 16

N=N=N How about: N N N This structure also satisfies the octet rule for each N atom Which structure is “correct”? Other possible structures for N 3 -

To determine which of the two structures for N 3 - is a more appropriate representation of the molecule we have to determine a quantity called the FORMAL CHARGE for each atom. The formal charge of an atom in a compound is the charge on the atom calculated on the assumption that the electrons in each bond are shared equally between the two atoms involved Formal charges assumes that all elements have the same electronegativity HClHCl Cl is more electronegative than H and hence the electrons associated with the covalent bond are “drawn” toward Cl

1 2 The formal charge of an atom in a compound is calculated as follows: formal charge = (group number) - (number of electrons in lone pairs) - (number of electrons in bonding pairs If H and Cl had the same electronegativity, the electron pair is equally shared by H and Cl, then H and Cl have the same number of electrons as before bonding Each is assigned a formal charge of 0 HCl 00

When two possible Lewis diagrams can be drawn for a molecule, the structure that assigns the lowest formal charge to the atoms is generally the best description of the bonding. When non-zero formal charges appear, Lewis structures that assign negative formal charges to the more electronegative atoms and positive ones to the more electropositive atoms are preferable. To decide which of the two structures for N 3 - is more appropriate calculate the formal charge on each atom for the two structures.

For (i) Formal charge on the left N = (1/2)(4) = -1 Formal charge on middle N = 5 - (1/2)(8) = +1 Formal charge on right N = (1/2)(4) = -1 N=N=NN N N (ii) (i) 1 2 formal charge = (group number) - (number of electrons in lone pairs) - (number of electrons in bonding pairs

For (ii) Formal charge on left N = (1/2)(2) = -2 Formal charge on middle N = 5 - (1/2)(8) = +1 Formal charge on right N = (1/2)(6) = 0 Sum of formal charge = = -1 (i) is the Lewis structure assigned to N 3 -, since (ii) assigns a formal charge of -2 on one of the N’s N=N=N (i) +1 N N N (ii) Note: sum of formal charge = charge on molecule This is a good check for a correct Lewis structure, and if the calculation for formal charge is correct.

So to the steps for drawing Lewis structures, add these steps 7) Determine the formal charge on each atom, and write it next to the atom. Check that the formal charges add to give the correct total charge of the molecule or molecular ion 8) If more than one Lewis structure is possible, then the structure that assigns the lowest formal charges to atoms is typically the best description of the bonding. When non-zero formal charges appear, Lewis structures that assign negative formal charges to the more electronegative atom and positive formal charges to the more electropositive atoms are preferable to the ones that do the reverse.

Keep in mind that formal charges do not represent real charges on atoms. The formal charge of an atom is the charge that an atom in a molecule would have if all atoms have the same electronegativity. Electronegativity differences between atoms determine the charge distribution in molecules.