Valence Electrons & Lewis Symbols

1. Valence Electrons & Lewis Symbols • Molecular compounds form when two atoms share valence electrons. • Electrons located in the outermost shell that can be transferred to or shared with another atom during the formation of ions or covalent bonds • Electrons over and above those found in the previous noble gas • The first step in determining the type of bonds formed between the atoms in a particular compound is to determine the number of valence electrons each atom has.

2. Valence Electrons & Lewis Symbols • For a main group element, • the number of valence electrons = group number for the element. • N (group 5A) has 5 valence electrons • Br (group 7A) has 7 valence electrons • We often use Lewis symbols (electron-dot symbols) as a way of depicting valence electrons and keeping track of them during the formation of chemical bonds.

3. Valence Electrons & Lewis Symbols • Components of a Lewis symbol: • Elemental symbol for the element • Dot for each valence electron • Dots are placed on all 4 sides of the elemental. • All four sides of the symbol are equivalent. • Up to 2 dots (e-) per side • Keep electrons unpaired until each side has 1 electron.

4. Lewis Symbols Example: Draw the Lewis symbol for each of the following atoms or ions. Nitrogen: Oxygen: Oxide ion:

5. Formation of Covalent Bonds • Atoms in a covalent bond are held together by the electrostatic attraction of the 2 nuclei for the concentration of negative charge between them. • Covalent bonds form when two atoms share electrons resulting in an increase in electron density between the two nuclei.

6. Formation of Covalent Bonds • Many atoms follow the octet rule when forming ions or forming covalent bonds. • Octet Rule:Atoms tend to gain, lose, or share electrons until surrounded by 8 valence electrons. • This gives the atom the same octet of electrons as all noble gases except helium have. • There are exceptions to the octet rule!

7. N N Formation of Covalent Bonds • In order to satisfy the octet rule, two atoms can form: • Single bonds • One pair of electrons are shared • Double bonds • Two pairs of electrons are shared • Triple bonds • Three pairs of electrons are shared

8. Lewis Structures • The covalent bonds that exist between atoms in a molecule or between atoms in a polyatomic ion are often depicted using Lewis structures. • Elemental symbol for each atom present • Covalent bond between atoms depicted using a solid line. • Unshared electron pairs are shown around the appropriate atom.

9. Lewis Structures • Steps for drawing Lewis structures for molecules or polyatomic ions: • Add up the valence electrons for all atoms & correct for the charge • For a cation (+), subtract 1 electron for each positive charge • NH4+: 5 + 4(1) -1 = 8 e- • For an anion (-), add 1 electron for each negative charge • CN-: 4 + 5 + 1 = 10 e-

10. Lewis Structures • Draw a skeleton structure showing which atoms are bonded to each other using a single bond (a straight line). • Many of the structures you draw will have a central atom surrounded by other (“outer”) atoms. • Use the atom other than hydrogen with the lowest electronegativity as the central atom • CCl4: C with 4 Cl attached • HCN: C with a N and a H attached

11. Atomic number Atomic weight (atomic mass) 5 10.81 B Element symbol electronegativity en = 2.04 boron Element name Lewis Structures • Electronegativity: • The ability of an atom that is part of a bond to attract electrons towards itself

12. Lewis Structures • Drawing skeletons (cont) • Oxyacids:the hydrogens in the formula are usually attached to one of the oxygens and NOT to the central atom • Some formulas are more complex, and I will have to give you “hints” to determine the skeleton.

13. Lewis Structures • Add pairs of electrons to the “outer” atoms first until each has an octet of electrons. • Exception: H can only have 2 electrons. • Place all leftover or unused electrons on the central atom. • Note: This may give the central atom more than 8 electrons in some cases.

14. Lewis Structures • Count the number of electrons around the central atom. • If there are fewer than 8 electrons on the central atom, use one or more unshared pairs of electrons from an outer atom to form a double or triple bond with the central atom. • Stop drawing double or triple bonds when you have satisfied the octet of the central atom.

15. Lewis Structures Example: Draw the Lewis structure for the following. Phosphate ion CHCl3

16. Lewis Structures Example: Draw the Lewis structure for the following. Nitrogen tribromide Xenon difluoride

17. Lewis Structures Example: Draw the Lewis structure for the following. Nitric acid HCO2H (Hint: C is the central atom. One H is bonded to O, and the other is bonded to C.

18. Lewis Structures • To determine which of two possible Lewis structures is more important (“best”), calculate the formal charge on each atom (other than H). • Formal charge:a calculated value that compares the # of valence electrons an atom has to the # of electrons assigned to it in a Lewis structure • FC = group # - nonbonding e- - 1/2 (bonding e-)

19. Lewis Structures • Use the formal charges to determine which Lewis structure is most important: • Greatest number of atoms with the smallest (closest to zero) formal charge • Negative charge on the more electronegative atom

20. Lewis Structures Example: Assign formal charges to each atom except hydrogen in the Lewis structures drawn for HCO2H. Which structure is the best representation for formic acid?

21. Lewis Structures • Common bonding patterns for “zero” formal charge: CNOHHalogens total bonds 4 3 2 1 1 lone pairs 0 1 2 0 3

22. Lewis Structures • Some substances cannot be adequately described by a single Lewis structure. • Ozone (O3): This single structure implies that the two O-O bonds should be different. They are not!

23. Lewis Structures • The structure for ozone can be depicted by two Lewis structures that differ only in the position of the electrons. • Lewis structures that are the same except for the position of the electrons are referred to asresonancestructures. • Ozone is a “hybrid” of the two resonance structures. • It has properties of both structures. • It does NOT “flip” back and forth.

24. Lewis Structures Example: Draw all possible resonance structures for: SCN- Nitrate ion:

25. Lewis Structures • Some compounds violate the octet rule: • Molecules with an odd # of electrons • NO (11 e-) • Molecules in which the central atom is electron deficient (less than 8 e-) • Molecules containing B as the central atom often are electron deficient.

26. Lewis Structures • Some compounds violate the octet rule (cont): • Molecules in which the central atom has more than an octet. • Central atoms from the 3rd period (or higher period #) can “expand their octet” and have more than 8 electrons.

27. Lewis Structures Example: Draw the Lewis structure for chlorine trifluoride.