Chapter 7 Bonding

What is a Bond? l A force that holds atoms together. l We will look at it in terms of energy. l Bond energy is the energy required to break a bond. l Why are compounds formed? l Because it gives the system the lowest energy.

Ionic Bonding l An atom with a low ionization energy reacts with an atom with high electron affinity. l The electron is transferred. l Opposite charges hold the atoms together.

Coulomb's Law l If charges are opposite, energy change is negative l exothermic l Same charge, positive E, requires energy to bring them together.

What about covalent compounds? l The electrons in each atom are attracted to the nucleus of the other. l The electrons repel each other, l The nuclei repel each other. l They reach a distance with the lowest possible energy. l The distance between is the bond length.

Covalent Bonding l Electrons are shared by atoms. l In polar bonds the electrons are not shared evenly. l One end is slightly positive, the other negative. Indicated using small delta 

Electronegativity l The ability of an atom to attract shared electrons to itself. l Tends to increase left to right. l decreases as you go down a group. l Noble gases aren’t part of it. l Difference in electronegativity between atoms tells us how polar bond is.

Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases

Dipole Moments l A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), l Also called a dipole moment. l Center of charge doesn’t have to be on an atom. l It line up in the presence of an electric field.

How It is drawn H - F ++ --

Which Molecules Have Them? l Any two atom molecule with a polar bond. l With three or more atoms there are two considerations: –There must be a polar bond. –Geometry can cancel it out.

Geometry and polarity l Three shapes will cancel them out. l Linear

Geometry and polarity l Three shapes will cancel them out. l Planar triangles 120º

Geometry and polarity l Three shapes will cancel them out. l Tetrahedral

Geometry and polarity l Others don’t cancel l Bent

Geometry and polarity l Others don’t cancel l Trigonal Pyramidal

Ions l Atoms tend to react to form noble gas configuration. l Metals lose electrons to form cations l Nonmetals can share electrons in covalent bonds. l Or they can gain electrons to form anions.

Ionic Compounds l Ions align themselves to maximize attractions between opposite charges, l and to minimize repulsion between like ions. l Can stabilize ions that would be unstable as a gas. l React to achieve noble gas configuration

Na(s) + ½F 2 (g)  NaF(s) First sublime NaNa(s)  Na(g)  H = 109 kJ/mol Ionize Na(g) Na(g)  Na + (g) + e -  H = 495 kJ/mol Break F-F Bond½F 2 (g)  F(g)  H = 77 kJ/mol Add electron to F F(g) + e -  F - (g)  H = -328 kJ/mol

Na(s) + ½F 2 (g)  NaF(s) Lattice energy Na(s) + ½F 2 (g)  NaF(s)  H = -928 kJ/mol l An ionic compound is considered any substance that conducts electricity when melted. l Also use the generic term salt.

The Covalent Bond l Due to the tendency of atoms to achieve the lowest energy state. l It takes 1652 kJ to dissociate a mole of CH 4 into its ions l Since each hydrogen is hooked to the carbon, we get the average energy = 413 kJ/mol

l CH 3 Cl has 3 C-H, and 1 C - Cl l the C-Cl bond is 339 kJ/mol l The bond is a method of explaining the energy change associated with forming molecules. l Bonds don’t exist in nature, but are useful as a model of behavior.

Covalent Bond Energies l We made some simplifications in describing the bond energy of CH 4 l Each C-H bond has a different energy. CH 4  CH 3 + H  H = 453 kJ/mol CH 3  CH 2 + H  H = 435 kJ/mol CH 2  CH + H  H = 425 kJ/mol CH  C + H  H = 339 kJ/mol

Localized Electron Model l A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. l Three Parts 1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals

Lewis Structure l Shows how the valence electrons are arranged. l One dot for each valence electron. l A stable compound has all its atoms with a noble gas configuration. l Hydrogen follows the duet rule, Be the quartet, B the sextet. l The rest follow the octet rule. l Bonding pair is the one between the symbols.

Rules l Sum the valence electrons. l Use a pair to form a bond between each pair of atoms. l Arrange the rest to fulfill the octet rule (except for H, Be and B). l H 2 O l A line can be used instead of a pair.

Practice l 1 st atom is commonly the central atom l POCl 3 l SO 4 -2 l SO 3 -2 l PO 4 -3 l SCl 2

Exceptions to the octet l SF 6 l Third row and larger elements can exceed the octet l Because they use 3d orbitals l I 3 -

Exceptions to the octet l When we must exceed the octet, extra electrons tend to go on central atom. l ClF 3 l XeO 3 l ICl 4 - l BeCl 2

Resonance l Sometimes there is more than one valid structure for an molecule or ion. l NO 3 - l Use double arrows to indicate it is the “average” of the structures. l It doesn’t switch between them. l NO 2 - l Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

Formal Charge l For molecules and polyatomic ions that exceed the octet there are several different structures. l Use charges on atoms to help decide which. l Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

Formal Charge l The difference between the number of valence electrons on the free atom and that assigned in the molecule. l We count half the electrons in each bond as “belonging” to the atom. l SO 4 -2 l Molecules try to achieve as low a formal charge as possible. l Negative formal charges should be on electronegative elements.

More Examples l XeO 3 l NO 4 -3 l SO 2 Cl 2

VSEPR l Lewis structures tell us how the atoms are connected to each other. l They don’t tell us anything about shape. l The shape of a molecule can greatly affect its properties. l Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

VSEPR l Molecules take a shape that puts electron pairs as far away from each other as possible. l Have to draw the Lewis structure to determine electron pairs. l Bonding or nonbonding lone pairs l Lone pair take more space.

VSEPR l The number of pairs determines –bond angles –underlying structure l The number of atoms determines –actual shape

VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar ° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octahedron

Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

If no central atom… l Can predict the geometry of each angle. l build it piece by piece.

How well does it work? l Predictions are almost always accurate. l But, like all simple models, it has exceptions.

Atomic Orbitals Don’t Work l to explain molecular geometry. l In methane, CH 4, the shape is tetrahedral. l The valence electrons of carbon should be two in s, and two in p. l the p orbitals would have to be at right angles. l The atomic orbitals change when making a molecule

l When two things come off. l One s and one p hybridize. l linear

sp hybridization l End up with two lobes 180º apart. l p orbitals are at right angles Makes room for two  bonds and two sigma bonds. l A triple bond or two double bonds.

sp 2 hybridization l When three things come off atom. l trigonal planar l 120º on  one  bond

sp 2 hybridization l C 2 H 4 l Double bond acts as one pair. l trigonal planar l Have to end up with three blended orbitals. l Use one s and two p orbitals to make sp 2 orbitals. l Leaves one p orbital perpendicular.

sp 3 Hybridization l We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. l We combine one s orbital and 3 p orbitals. l sp 3 hybridization has tetrahedral geometry.

How we get to hybridization l We know the geometry from experimentation l We know the orbitals of the atom l hybridizing atomic orbitals can explain the geometry. l So if the geometry requires a tetrahedral shape, it is sp 3 hybridized l This includes bent and trigonal pyramidal molecules because one of the sp 3 lobes holds the lone pair.

sp 3 d hybridization l The model predicts that we must use the d orbitals. l sp 3 d hybridization l sp 3 d hybridization, ex: PCl 5 l Trigonal bipyrimidal can only  bond. can’t  bond. l basic shape for five bonds.

sp 3 d 2 l gets us to six bonds around l octahedral

Two types of Bonds l Sigma bonds from overlap of orbitals. l Between the atoms. Pi bond (  bond) above and below atoms l Between adjacent p orbitals. l The two bonds of a double bond.

CO 2 C can make two  and two  O can make one  and one  COO