Chapter 6 NOR AKMALAZURA JANI CHM 138 BASIC CHEMISTRY

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding. 1A 1ns 1 2A 2ns 2 3A 3ns 2 np 1 4A 4ns 2 np 2 5A 5ns 2 np 3 6A 6ns 2 np 4 7A 7ns 2 np 5 Group# of valence e - e - configuration

Lewis Dot Symbols for the Representative Elements & Noble Gases Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element.

Li + F Li + F - The Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 Ionic bond: the electrostatic force that holds ions together in an ionic compound. Atoms of the elements with low ionization energies tend to form cation – alkali metals and alkaline earth metals Atoms of the elements with high electron affinities tend to form anion – halogens and oxygen (LiF)

Li Li + + e - e - + FF - F - Li + + Li + F -

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond

A Lewis structure: - a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots an individual atoms. Only valence electrons are shown. The formation of the molecules illustrates the octet rule. - Octet rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons.

8e - H H O ++ O HH O HHor 2e - Lewis structure of water Double bond – two atoms share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds Triple bond – two atoms share three pairs of electrons N N 8e - N N triple bond or

Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

COMPARISON OF GENERAL PROPERTIES OF IONIC COMPOUND AND COVALENT COMPOUND IONIC COMPOUND COVALENT COMPOUND Solid at room temperature, high melting points Gases, liquids, low melting solids Soluble in water Insoluble in water Aqueous solution conduct electricity Aqueous solution do not conduct electricity Strong electrolytes Nonelectrolytes

1.Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. 2.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 3.Complete an octet for all atoms except hydrogen 4.If structure contains too many electrons, form double and triple bonds on central atom as needed. Writing Lewis Structures

Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – N is less electronegative than F, put N in center FNF F Step 2 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e - 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

Write the Lewis structure for: i) NO - 2 i) NO - 2 ii) CS 2 ii) CS 2 iii) SO 3 iii) SO 3

FORMAL CHARGE An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number electron assigned to atom The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

HCOH Examples: H CO H formal charge on C = 0 formal charge on O = 0 formal charge on C = -1 formal charge on O = + 1

Formal Charge and Lewis Structures 1.For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. 2.Lewis structures with large formal charges are less plausible than those with small formal charges. 3.Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH 2 O? HCOH +1 H CO H 00

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. OOO + - OOO + - OCO O -- OCO O - - OCO O - - What are the resonance structures of the carbonate (CO 3 2- ) ion? Resonance Structure

Exceptions to the Octet Rule The Incomplete Octet HHBe Be – 2e - 2H – 2x1e - 4e - BeH 2 BF 3 B – 3e - 3F – 3x7e - 24e - FBF F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24

Odd-Electron Molecules N – 5e - O – 6e - 11e - NO N O The Expanded Octet (central atom with principal quantum number n > 2) SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = lone pairs (18x2) = 36 Total = 48 Exceptions to the Octet Rule

Dative Covalent Bond / Coordinate Covalent Bond A covalent bond in which one of the atoms donates both electrons. Examples: - NH 4 +, NH 3 AlCl 3, NH 3 BF 3

Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. 1.Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. 2.Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. 3.Covalent bonds are formed by: a.Overlap of hybrid orbitals with atomic orbitals b.Overlap of hybrid orbitals with other hybrid orbitals Hybridization

Formation of sp 3 Hybrid Orbitals

Formation of Covalent Bonds in CH 4

sp 3 -Hybridized N Atom in NH 3

Formation of sp Hybrid Orbitals

Formation of sp 2 Hybrid Orbitals

# of Lone Pairs + # of Bonded Atoms HybridizationExamples sp sp 2 sp 3 sp 3 d sp 3 d 2 BeCl 2 BF 3 CH 4, NH 3, H 2 O PCl 5 SF 6 How to predict the hybridization of the central atom? 1.Draw the Lewis structure of the molecule. 2.Count the number of lone pairs AND the number of atoms bonded to the central atom

Sigma bond (  ) – covalent bonds formed by orbitals overlapping end – to-end, with the electron density between the nuclei of the bonding atoms Pi bond (  ) – a covalent bond formed by sideways overlapping orbitals with electron density concentrated above and below plane of nuclei of the bonding atoms Bonding in Ethylene, C 2 H 4

Another View of  Bonding in Ethylene, C 2 H 4

Bonding in Acetylene, C 2 H 2

Describe the bonding in CH 2 O C H O H C – 3 bonded atoms, 0 lone pairs C – sp 2

Sigma (s) and Pi Bonds (p) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many  and  bonds are in the acetic acid (vinegar) molecule CH 3 COOH? C H H CH O OH  bonds = = 7  bonds = 1

Intermolecular Forces Intermolecular forces: attractive forces between molecules. Van der Waals forces: - the attractive or repulsive force between molecules due to covalent bonds or to the electrostatic interaction of ions with one another or with neutral molecules. The term includes: - permanent dipole–permanent dipole forces - instantaneous induced dipole-induced dipole (London dispersion force). Examples: interaction between H 2, Cl 2, F 2, CH 4

Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B A H … A or A & B are N, O, or F

Why is the hydrogen bond considered a “special” dipole-dipole interaction? Decreasing molar mass Decreasing boiling point