Chapter 6 Chemical Bonding Bonding Theory and Electronegativity

Chapter 6 Chemical Bonding

Bonding Theory and Electronegativity • Atoms want to achieve noble gas configuration 8 valence e • Some elements have stronger attractions to ewhen bonded • ELECTRONEGATIVITY (EN) – Relative attraction an atom has for shared electrons in a covalent bond – Unit- paulings – Arbitrary number used for comparison purposes • F is 4. 0, Cs/Fr 0. 7

– Increase from left to right in a period – Decrease from top to bottom in a group – Nonmetals higher than metals – Metals on left side and nonmetals on right side most reactive (alkali metals and halogens) – Electrons attracted to the higher EN element

Determining Bond Types Generally two atoms of same element or same EN value Nonmetal bonded to nonmetal with different EN values Metal bonded to nonmetal

Why do atoms bond? • Noble gas atoms are unreactive because their electron configurations are especially stable. – Outer energy level has eight electrons. (except He) • Other atoms fill their outermost energy levels by bonding. • octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level.

Ionic Bonding • Bonds share electrons unevenly – Depends on difference between EN of atoms • Metals tend to lose e- to get stable • Nonmetals tend to gain eto get stable • Lewis diagrams help to show e- move

Covalent Bonding • Covalent bonding happens with nonmetals, forming MOLECULES • Atoms tend to fill their outer energy levels. OCTET RULE • Unpaired e- are available to bond (SHARE) • When two nonmetals combine, a covalent bond may form where electrons are shared • Explained by simultaneous attraction of the shared pair by both nuclei • Number of covalent bonds each atom can form is limited by how many are needed for octet. – C needs 4, O needs 2, etc.

Lewis Molecular Theory- Lewis Diagrams and Structural Formulas • Lewis Diagram of atoms drawn normally using valence e- F Four bonding pairs Three lone pairs, one bonding pair

• Lewis Diagram of Molecule shares the bonding electrons between atoms • Structural Formuladrop lone pairs and put a dash for shared pairs

Writing Lewis Diagrams for Molecules • Steps – Count all valence e– Draw skeleton structure – Put a pair of e- between all atoms to show a covalent bond • All should have 8 (exc H which has 2) – Distribute lone pairs around atoms (exc H) – If an atom needs more e- then move pairs between atoms to get 8 • Most multiple bonds are in C, N, and O • Double bond (share two), triple bond (share three)

• A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms. – Double bonds are often found in molecules containing carbon, nitrogen, and oxygen. • A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes

• A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms. • example 1—diatomic nitrogen: • example 2—ethyne, C 2 H 2:

Stereochemistry- VSEPR Theory • All molecules have 3 D shape • Stereochemistry- study of shapes of molecules • VSEPR theory – Valence Shell Electron Pair Repulsion – Helps to understand predict molecular geometry (from Lewis Dot diagrams) – Developed by Gillespie and Nyholm in 1956 -57 – Rules based on the idea that the arrangement in space of the covalent bonds formed by an atom depends on the arrangement of valence e • e- try to push each other far away while still bonding to central atom • e- pair repulsion

• Restricted VSEPR rules – Valence e- pairs (both shared and lone) arrange themselves around the central atom in a molecule in such a way as to minimize repulsion (as far away from each other as possible) – When predicting molecular geometry, double and triple bonds act like single bonds – Lone pairs of e- occupy more space than bonding e-

• Steps to draw VSEPR molecules – Draw Lewis Diagram – Determine the central atom (lowest EN) – Count the number of bonding and lone pairs surrounding the central atom • Multiple bonds count as one pair – Shape molecule in order to minimize repulsion • Find on VSEPR chart

EXAMPLES • Water, H 2 O – 2 bond pairs – 2 lone pairs – The molecular geometry is BENT.

– As a Lewis Dot diagram: – Has two electron pairs - two sets of double bonds – two bonding pairs – Creates a LINEAR shape according to VESPR

Polarity of Covalent Bonds • Polarity of a molecule is determined by the behavior of bonding pairs and the difference in EN • Nonpolar Covalent – The bonding electron pair is shared equally and is uniformly distributed between the nuclei of two bonded atoms • Only atoms of the same element or with a difference in EN of 0. 0 -0. 4 • EX- H 2, F 2, I 2, N 2, CS 2

• Polar Covalent – The bonding electrons are unequally shared and unevenly distributed between the nuclei of two unbonded atoms – Difference in EN makes one atom hold e- more strongly – HCl is POLAR because it has a positive end a negative end. (difference in electronegativity) – Cl has a greater share in bonding electrons than does H. – Cl has slight negative charge (-d) and H has slight positive charge (+ d)