Chapter 11 Chemical Bonding

11.1 Characteristics of Chemical Bonds Chemical bond – the force that holds two or more atoms together and makes them function as a unit Bond energy – energy required to break a given chemical bond

Types of Chemical Bonds Ionic Bond – attraction between oppositely charged ions Electrons lost/gained

Covalent Bond – type of bond in which atoms share electrons

Polar covalent bond – a covalent bond in which the electrons are not shared equally More electronegative atom pulls electrons closer

Electronegativity Electronegativity - The tendency of an atom in a molecule to attract shared electrons to itself Increases across a period Decreases down a group

Rank the following in order of increasing electronegativity (without looking at the values): Al, Cl, S, F, Na

Polarity of a bond depends on the difference between the electronegativity values of the atoms

Rank the following in order of increasing polarity: H-H, O-H, Cl-H, S-H, F-H

Bond Polarity and Dipole Moments A dipole moment results when a polar molecule has a center of positive charge separate from a center of negative charge

Water has a dipole moment which is the reason you are sitting here RIGHT NOW!!!

Liquid at room temperature “universal solvent” Like dissolves like Liquid at room temperature

11.2 Characteristics of Ions and Ionic Compounds

Ions form to achieve a noble gas electron configuration Octet rule – atoms will lose, gain, or share electrons to acquire a full set of 8 valence electrons Metals nonmetals

Neutral atom Na 1s22s22p63s1 B 1s22s22p1 P 1s22s22p63s23p3 F 1s22s22p5

Ion Na1+ 1s22s22p6 (Ne) B3+ 1s2 (He) P3- 1s22s22p63s23p6 (Ar) F1-

1 = s1 (1 ve, easily lost, ion formed = ) 13 = s2p1 14 = s2p2 15 = s2p3 16 = s2p4 17 = s2p5 18 = s2p6

Ionic Bonding and Structures of Ionic Compounds LiF = formula unit Actual compound contains a lot of ions packed together to maximize attraction Every + ion surrounded by – ions Every – ion surrounded by + ions

Ionic compounds containing polyatomic ions Behave the same way as monatomic ions Covalent bonds hold polyatomic ion together so it behaves as a unit

11.3 Lewis Structures Lewis structure – representation of a molecule or polyatomic ion showing how electrons are arranged among the atoms Only include valence electrons

Bonding pair – a pair of electrons shared between 2 atoms Lone pair – electron pairs in a Lewis structure that are not involved in bonding

If atoms need 1 electron, it will usually form 1 covalent bond. H and Halogens typically only form one bond If atoms need 2 electrons, it will usually form 2 covalent bonds. If atoms needs 3 electrons, it will usually form 3 covalent bonds.

Steps for writing Lewis Structures: Obtain the sum of the valence electrons from all of the atoms Use one pair of electrons to form a bond between each pair of bound atoms Arrange the remaining electrons to satisfy the octet rule Exception:

Single Covalent Bonds One pair of electrons is shared Cl2 HF

Double Covalent Bond Two pairs of electrons are shared O2 Typically C, N, O, S O2

Triple covalent bond Three pairs of electrons are shared N2 Typically C or N N2

Write the Lewis structures for the following: NH3 CO2

C2H4

Lewis Structures of Polyatomic Ions Number of dots in the molecule must take the ions charge into account NH41+

PO43-

Resonance Structures A molecules shows resonance when more than one Lewis structure can be drawn for the molecule

O3 CO32-

Exceptions to the octet rule Suboctets – stable configurations with fewer than 8 electrons around an atom Boron typically forms a suboctet BH3

Expanded octets – central atoms contain more than eight valence electrons Phosphorous and sulfur can form expanded octets PCl5

12.4 Structures of Molecules VSEPR model V alence S hell E lectron P air R epulsion The shape of a molecule is determined by minimizing repulsion between lone pairs

Bond angle VSEPR forces cause atoms in a molecule to be positioned at fixed angles relative to one another

Linear – 180 degrees Binary molecules or molecules with 2 atoms bonded to the central atom and no lone pairs on the central atom CO2

Bent – 104.5 degrees Two atoms bonded to central atom, two lone pairs on central atom H2O

Trigonal Planar – 120 degrees Three atoms bonded to central atom, no lone pairs on central atom BH3

Tetrahedral – 109.5 degrees Pyramid shaped 4 atoms bonded to central atom, no lone pair on central atom CH4

Trigonal Pyramidal – 107.3 degrees A triangle that is not flat Three atoms bonded to central atom, one lone pair on central atom NH3

Polar and Nonpolar Molecules Polar bond = electrons shared unequally (dipole) More electronegative atom pulls electrons closer to it and therefore has a slightly negative charge Less electronegative atom has a slightly positive charge

Polar molecules always have: At least one polar covalent bond Asymmetric geometry Lone pairs on the center atom Different atoms bonded to the center atom

HCl H2O

Not every molecule with polar bonds is polar!!! CH4 CO2

Polar Bond vs. Polar Molecule Polar bond = electrons shared unequally between molecules Polar molecule = entire molecule has different partial charges on opposite sides of the molecule Depends on shape

CCl4 NH3 H20 Linear vs. bent