Ch. 10: Chemical BondingCh. 10: Chemical Bonding

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

I. Chapter OutlineI. Chapter Outline

I. Introduction

II. Lewis Dot Structures

III. Ionic Compounds

IV. Covalent Compounds

V. Shapes of Molecules

VI. Polar Bonds and Polar Molecules

I. IntroductionI. Introduction

• Bonding theories allow prediction of how atoms form compounds and the shapes they will take.

• The 3-D shape of a molecule determines many of its physical properties.

• With the advent of faster computers, simulations allow screening of drug candidates.

I. HIV-protease InhibitorsI. HIV-protease Inhibitors

I. Lewis TheoryI. Lewis Theory

• Lewis theory is simple to apply, but amazingly powerful.

• It can be used to go from a formula of a compound to its 3-D structure.

• Lewis theory centers on how valence electrons are used by atoms to form compounds.

II. Lewis Dot StructuresII. Lewis Dot Structures

• Valence e-’s are the most important e-’s in bonding.

• Lewis dot structures are a way to depict the valence e-’s of atoms.

• Lewis dot symbols have two parts:1) element symbol: represents nucleus and

core e-

2) dots around symbol: represent valence e-’s

II. Origin of Dot StructuresII. Origin of Dot Structures

• Oxygen has 6 valence electrons, so it’s dot structure will have 6 dots.

II. Lewis Dot StructuresII. Lewis Dot Structures

• The number of valence e- is given by the element’s group number!!

II. The Central IdeaII. The Central Idea

• Lewis realized that noble gases all had 8 valence e-’s (except He).

• He reasoned that having 8 valence e-’s (or 2 for H and He) leads to stability, or being unreactive.

• These special configurations of valence e-’s are known as an octet or a duet.

II. The Octet RuleII. The Octet Rule

• In Lewis theory, atoms bond in order to obtain eight valence electrons.

• The octet rule states that in chemical bonding, atoms transfer or share electrons to obtain outer shells with eight electrons.

• The octet rule is a main-group concept, generally applying to all except H and Li.

II. Sample ProblemII. Sample Problem

• Draw Lewis dot structures for Ca, Ge, Se, and Br.

III. Transfer of ElectronsIII. Transfer of Electrons

• When metals bond with nonmetals, the metal transfers electrons to the nonmetal to form an ionic compound.

• The positive charge of the metal cation and the negative charge of the nonmetal anion holds the ionic compound together.

• We can show this with Lewis theory.

III. Formation of Potassium III. Formation of Potassium ChlorideChloride

III. Formation of Sodium III. Formation of Sodium SulfideSulfide

III. Sample ProblemIII. Sample Problem

• Use Lewis dot structures to depict the compound that forms between: magnesium and nitrogen aluminum and bromine

IV. Sharing ElectronsIV. Sharing Electrons

• When nonmetals bond with other nonmetals, a covalent compound is formed.

• Electrons are shared in covalent compounds in order to achieve an octet.

• Electrons that appear in the space between atoms count towards the octet of both atoms.

IV. Lewis Structure for WaterIV. Lewis Structure for Water

IV. Octet/Duets SatisfiedIV. Octet/Duets Satisfied

IV. Bonding vs. Lone PairsIV. Bonding vs. Lone Pairs

• Electrons shared between atoms are bonding pair electrons.

• Electrons that are only on one atom are lone pair or nonbonding electrons.

IV. Lines as BondsIV. Lines as Bonds

• Generally, bonding pair electrons are represented by lines.

• Note that a line equals two electrons.

IV. Why Some Elements Exist IV. Why Some Elements Exist as Diatomicsas Diatomics

• If two hydrogens combine, they can satisfy their duet.

• If two chlorines combine, they can satisfy their octet.

IV. DiatomicsIV. Diatomics

IV. Higher Order BondsIV. Higher Order Bonds

• In some compounds, atoms need to share more than one electron pair to reach an octet.

• Double bond – atoms share 4 electrons

• Triple bond – atoms share 6 electrons

• However, higher order bonds are a last resort used by atoms to reach an octet!

IV. Diatomic OxygenIV. Diatomic Oxygen

• Oxygen has 6 valence electrons.

• Sharing one pair satisfies the octet of one oxygen atom, but not the other.

• Since there are no more electrons that can be used, higher order bonds must be made.

IV. Double Bond FormationIV. Double Bond Formation

IV. Triple BondsIV. Triple Bonds

• Sometimes six electrons need to be shared by two atoms.

• An example is diatomic nitrogen.

IV. Single, Double, Triple BondsIV. Single, Double, Triple Bonds

• Higher order bonds mean more stability and shorter internuclear distances.

Bond Type Internuclear Distance (pm)

Single 148

Double 121

Triple 110

IV. Steps for Drawing Lewis StructuresIV. Steps for Drawing Lewis Structures1) Determine total # of valence e-. (Cation,

subtract e-’s for charge; anion, add e-’s for charge).

2) Place atom w/ lower Group # (lower electronegativity) as the central atom.

3) Attach other atoms to central atom with single bonds.

4) Fill octet of outer atoms. (Why?)5) Count # of e- used so far. Place

remaining e- on central atom in pairs.6) If necessary, form higher order bonds to

satisfy octet rule of central atom.7) Allow expanded octet for central atoms

from Period 3 or lower.

IV. Sample ProblemIV. Sample Problem

• Draw correct Lewis structures for NF3, CO2, SeCl2, CCl4, and H2CO.

IV. Exceptions to the Octet RuleIV. Exceptions to the Octet Rule

• Lewis theory is too simple to cover all bonding possibilities.

• Some exceptions exist to the octet rule: e- deficient atoms like Be and B, e.g. BeCl2

and BF3.

Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO2.

Expanded valence – when d orbitals are used to accommodate more than an octet.

V. VSEPR TheoryV. VSEPR Theory

• From a correct Lewis structure, we can get to the 3-D shape using this theory.

• VSEPR stands for valence shell electron pair repulsion.

• The theory is based on the idea that electron groups (lone pairs, single bonds, or multiple bonds) repel each other.

V. Linear GeometryV. Linear Geometry

• CO2 has two electron groups.

• The two double bonds try to get as far away from each other as possible.

V. Trigonal PlanarV. Trigonal Planar

• Formaldehyde has three electron groups around the central atom.

• They form 120° angles to get away from each other.

V. TetrahedralV. Tetrahedral

• Methane has four electron groups.

• A bond angle of 109.5° keeps them furthest apart.

V. Tetrahedral-Based ShapesV. Tetrahedral-Based Shapes

• Other shapes are based on tetrahedral with bonding groups being replaced by lone pair electrons.

• The electron geometry is the shape based on all electron types (bonding and lone pair).

• The molecular geometry is the shape based on just atoms.

V. AmmoniaV. Ammonia

• The central atom has three bonds and one lone pair (4 electron groups).

• EG = tetrahedral

• When drawing MG, lone pairs are left off!

V. Trigonal PyramidalV. Trigonal Pyramidal

• Ammonia has a trigonal pyramidal molecular geometry.

V. WaterV. Water

• The central atom has two bonds and two lone pairs (4 electron groups).

• EG = tetrahedral

• Again, lone pairs omitted when drawing MG!

V. BentV. Bent

• Water has a bent molecular geometry.

V. Geometry SummaryV. Geometry SummaryElectronGroups

BondingGroups

LonePairs

ElectronGeometry

Ideal Angle

MolecularGeometry

2 2 0 Linear 180° Linear

3 3 0 Trigonalplanar

120° Trigonalplanar

3 2 1 Trigonalplanar

120° Bent

4 4 0 Tetrahedral

109.5° Tetrahedral

4 3 1 Tetrahedral

109.5° Trigonal pyramidal

4 2 2 Tetrahedral

109.5° Bent

V. Drawing w/ PerspectiveV. Drawing w/ Perspective

• We use the conventions below to depict a 3-D object on a 2-D surface.

V. Practice ProblemV. Practice Problem

• Draw the molecular shapes for ClO2-,

BF3, and NF3. Indicate the name of the molecular and electronic geometries for each as well.

VI. Sharing ElectronsVI. Sharing Electrons

• It is not reasonable to assume that all atoms will share electrons equally.

• Some atoms will pull electrons closer to them.

• Electronegativity is the ability of an element to pull electrons in a covalent bond closer.

VI. Effect of ElectronegativityVI. Effect of Electronegativity

• O is more electronegative than H, so in an O-H bond, the bonding electrons are more likely found around the O.

VI. Dipole MomentVI. Dipole Moment• The unequal sharing leads to a partial

charge separation in the bond called a dipole moment.

• Covalent bonds with a dipole moment are polar covalent bonds.

• The greater the difference in electronegativity, the greater the polarity.

VI. Electronegativity ValuesVI. Electronegativity Values

VI. Bonding TypeVI. Bonding Type

• Electronegativity difference can be used to determine the type of bond.

VI. Sample ProblemVI. Sample Problem

• Calculate the difference in electronegativity for the following pairs of atoms and determine of the bond between them is pure covalent, polar covalent, or ionic. I and I Cs and Br P and O

VI. Polar MoleculesVI. Polar Molecules

• Just because a molecule has polar bonds doesn’t mean that the molecule is polar overall.

• The degree of polarity in each bond and the orientation of those polar bonds determines whether the molecule is polar.

VI. Polarity VectorsVI. Polarity Vectors

• We can use vectors to represent dipole moments and analyze the resultant.