Chemistry-ch05_chemical Bonding And Molecular Structure

5

Chemical bonding and molecular structure

5.1 Fundamentals of bonding • There are three types of interactions within a molecule: – Electrons and nuclei attract one another – Electrons repel each other – Nuclei repel each other A

• In any molecule, these three interactions are balanced to give the molecule its greatest possible stability

5.1 Fundamentals of bonding • This balance is achieved when the electron density is situated between the nuclei of bonded atoms • This shared electron density is called a covalent bond • The attractive energy between nuclei and electrons exceeds the repulsive energy arising from nuclei-nuclei and electron-electron interactions

5.1 Fundamentals of bonding • The hydrogen molecule – These concepts are demonstrated below for the simplest stable neutral molecule, molecular hydrogen

5.1 Fundamentals of bonding • Bond length and bond energy

5.1 Fundamentals of bonding – Bond length is the separation distance at which the molecule has the maximum energetic advantage over the separated atoms – Bond energy is the energy required to break the bond (kJ mol-1) – Bond energy is always positive – Each different chemical bond has a characteristic bond length and energy

5.1 Fundamentals of bonding • Other diatomic molecule: F2 – Even when the atoms within a molecule contain many electrons, bond formation is still considered as the sharing of only two electrons – The resulting bond is a sigma bond

5.1 Fundamentals of bonding • Unequal electron sharing – In a chemical bond between any two identical atoms, the nuclei share the bonding electrons equally – When it is two different atoms, unequal attractive forces lead to an unsymmetrical distribution of the bonding electrons – This results in a polar covalent bond

5.1 Fundamentals of bonding – Electronegativity gives a numerical value of how strongly an atom attracts the electrons in a chemical bond

5.2 Ionic bonding • Compounds formed between elements of very different electronegativities are ionic in character • Most ionic compounds are solids with high melting points • Ionic compounds are held together in 3D arrangements by the attractive forces between oppositely charged ions

5.2 Ionic bonding

5.2 Ionic bonding • The lattice energy of a compound is dependent on the charges of the ions

5.3 Lewis structures • The conventions – Each atom is represented by its elemental symbol – Only the valence electrons appear in a Lewis structure – A line represents one pair of electrons that is shared between two atoms – Dots represent the nonbonding electrons on that atom

5.3 Lewis structures

5.3 Lewis structures – Procedure for drawing Lewis structures: • Count the valence electrons • Assemble the bonding framework using single bonds • Place three nonbonding pairs of electrons on each outer atom except H • Assign the remaining valence electrons to inner atoms • Minimise formal charges on all atoms

5.3 Lewis structures • Building Lewis structures – Lewis structures by themselves do not imply any particular geometry for a molecule or ion – A set of four electrons associated with an atom is often called an octet – Larger atoms can accommodate more than eight electrons, eg sulfur atom – The nonbonding pairs are called lone pairs

5.3 Lewis structures • Resonance structures – Are composites of equivalent Lewis structures – They are connected by double-headed arrows to emphasise that a complete depiction requires all of them

5.3 Lewis structures – Resonance structures differ only in the position of the electrons, not atoms – Resonance structures are not always equivalent – Resonance structures that are not equivalent occur when electrons are shifted between atoms of different elements

5.4 Valence shell electron pair repulsion (VSEPR) theory

• The VSEPR theory considers that molecular shape is determined primarily by the repulsions between pairs of electrons in the molecule • To minimise these repulsions, electron pairs around an inner atom within a molecule will be situated as far apart as possible

5.4 Valence shell electron pair repulsion (VSEPR) theory

• The procedure:

– Draw the Lewis structure of the molecule – Count the number of sets of bonding pairs and lone pairs of electrons – Use the following table:

5.4 Valence shell electron pair repulsion (VSEPR) theory

• Two sets of electron pairs: linear geometry – Example: CO2

• Have two sets of electron pairs that need to be situated as far away from each other as possible • This is satisfied by placing the two sets in a linear arrangement and results in a linear shape

5.4 Valence shell electron pair repulsion (VSEPR) theory

• Three sets of electron pairs: trigonal planar geometry – Example: BF3

• Have three sets of electron pairs around an inner atom • In trigonal planar, the sets are oriented at 120° to each other and are coplanar

5.4 Valence shell electron pair repulsion (VSEPR) theory

• Four sets of electron pairs: tetrahedral geometry – Example: CH4

• Four sets of equivalent electron pairs • The optimal geometry is tetrahedral, with H atoms situated at the four vertices of a tetrahedron

5.4 Valence shell electron pair repulsion (VSEPR) theory

– Example: NH3

• Total of four sets of electron pairs, three bonding and one lone pair • The single lone pair on the N atom takes up more space than each of the three bonding pairs • Adopts a trigonal pyramidal shape

5.4 Valence shell electron pair repulsion (VSEPR) theory

– Example: H2O

• Four sets of electron pairs, two bonding and two lone pairs • The idealised tetrahedral geometry is significantly distorted • Water adopts a bent shape

5.4 Valence shell electron pair repulsion (VSEPR) theory

• Five sets of electron pairs: trigonal bipyramidal geometry – Example: PCl5 • Five sets of equivalent electron pairs

5.4 Valence shell electron pair repulsion (VSEPR) theory

5.4 Valence shell electron pair repulsion (VSEPR) theory

• Six sets of electron pairs: octahedral geometry – Example: SF6

• Six sets of electron pairs around the S atom • An octahedral geometry of the six sets places them as far apart as possible

5.4 Valence shell electron pair repulsion (VSEPR) theory • All six positions around an octahedron are equivalent

5.4 Valence shell electron pair repulsion (VSEPR) theory

5.5 Properties of covalent bonds • Dipole moments – A molecule with asymmetrical distribution of electron density is said to have a dipole moment (μ)

5.5 Properties of covalent bonds – Dipole moments depend on bond polarities

– They also depend on molecular shape

5.5 Properties of covalent bonds • Bond length – Bond length of a covalent bond is the nuclear separation distance at which the molecule is most stable – At this distance, attractive interactions are maximised relative to repulsive interactions

5.5 Properti es of covalent bonds

5.5 Properties of covalent bonds

5.5 Properties of covalent bonds • The following factors influence bond length: – The smaller the atoms, the shorter the bond – The more electrons in a bond, the shorter the bond – The larger the electronegativity difference between the bonded atoms, the shorter the bond

5.5 Properties of covalent bonds • Bond energy – The amount of energy that must be supplied to break a chemical bond – Three consistent trends: • Bond energies increase as more electrons are shared between the atoms • Bond energies increase at the electronegativity difference between bonded atoms increases • Bond energies decrease as bonds become larger

Summary of molecular shapes

5.6 Valence bond theory • Two ways to think about bonding: – Valence bond theory • Assumes that electrons are either localised in bonds between two atoms or localised on a single atom, usually in pairs • Considers two types of orbitals

– Molecular orbital theory • Uses delocalised bonds • Requires a more complicated analysis, but explains chemical properties that localised bonds cannot

5.6 Valence bond theory • Orbital overlap – Bonding orbitals are created by combining atomic orbitals – Insert Figure 5.20, p178

5.6 Valence bond theory – Orbital overlap occurs when two orbitals of the same phase are superimposed and result in a new orbital

5.6 Valence bond theory • Conventions of the orbital overlap model – Each electron in a molecule is assigned to a specific orbital – No two electrons in a molecule have identical descriptions – The electrons in molecules obey the aufbau principle – Only the valence orbitals are needed to describe bonding

5.6 Valence bond theory • Hybridisation of atomic orbitals – Hybrid orbitals are combinations of atomic orbitals – The process by which we combine them is called hybridisation – Example: methane

5.6 Valence bond theory – Methane: sp3 hybrid orbitals • The valence orbitals of interest are the one 2s and three 2p orbitals • These four atomic orbitals are ‘mixed’ to form four new hybrid orbitals

5.6 Valence bond theory

• The total energy of the four atomic orbitals must equal the total energy of the four hybrid orbitals • The four sp3 orbitals position themselves so that their electrons undergo the minimum repulsion

5.6 Valence bond theory

• sp3 hybridisation is not limited to C atoms. An atom which exhibits a tetrahedral arrangement of electron pairs can be considered to be sp3 hybridised

5.6 Valence bond theory – sp2 hybrid orbitals • Mixture of the 2s orbital and two 2p orbitals

5.6 Valence bond theory • The three sp2 orbitals are coplanar

5.6 Valence bond theory – sp hybrid orbitals • Hybridisation of a 2s orbital with a 2p orbital

5.6 Valence bond theory

5.6 Valence bond theory • Multiple bonds – Bonding in ethene • Charges can be minimised by forming a double bond between the carbon atoms, leaving both carbons with an octet of electrons • The carbon atoms are sp2 hybridised • Trigonal planar geometry

5.6 Valence bond theory

5.6 Valence bond theory • The valence bond description of any double bond can be constructed using the following procedure: – Determine the Lewis structure – Use the Lewis structure to determine the type of hybridisation – Construct the σ bond framework – Add the π bonds

5.6 Valence bond theory – Bonding in ethyne • The carbon atoms are joined by a triple bond • sp hybrid orbitals describe the bonding

5.6 Valence bond theory

• Each triple bond has one C-C σ bond and two C-C π bonds, which is always the case in any alkyne

5.7 Molecular orbital theory: diatomic molecules • Molecular orbital theory can predict and explain molecular properties • Differs fundamentally from valence bond theory • Electrons within a molecule are not localised, instead they occupy molecular orbitals (MOs)

5.7 Molecular orbital theory: diatomic molecules • Molecular orbitals of H2 and He2 – The overlap of N atomic orbitals will always lead to the formation of N molecular orbitals • In-phase overlap is constructive and results in the formation of a molecular orbital with a large amplitude • Out-of-phase overlap is destructive and gives a molecular orbital with zero amplitude • The region of zero amplitude is called a node

5.7 Molecular orbital theory: diatomic molecules

5.7 Molecular orbital theory: diatomic molecules – Constructive overlap gives a bonding molecular orbital • Electron density is maximised between the two nuclei

– Destructive overlap gives an antibonding molecular orbital • Electron density is minimised between the two nuclei and is zero at the node

– Molecular orbital diagrams show the relative energies of atomic and molecular orbitals

5.7 Molecular orbital theory: diatomic molecules • The bonding orbital is of lower energy than the orbitals from which it was formed • The antibonding orbital is of higher energy

5.7 Molecular orbital theory: diatomic molecules – Electrons of the molecule are assigned to orbitals of successively higher energy according to the aufbau principle and Hund’s rule • Electrons occupy the lowest energy orbitals first. • Atoms are most stable with the highest number of unpaired electrons

5.7 Molecular orbital theory: diatomic molecules

– Bond order represents the net amount of bonding between two atoms BO =

1 ( # electrons in bonding MOs − # electrons in antibonding MOs ) 2

– A bond order of 1 corresponds to a single bond

Molecular orbitals of O2

5.7 Molecular orbital theory: diatomic molecules – Rules: • The bonding and antibonding σs orbitals are of lower energy than any of the six molecular orbitals derived from the 2p orbitals • The two π bonding orbitals are degenerate, because the corresponding atomic p orbitals from which they are derived are degenerate • The antibonding orbitals formed from the atomic 2p orbitals are highest in energy, with the σ*p orbital higher than the π*

5.7 Molecular orbital theory: diatomic molecules

5.7 Molecular orbital theory: diatomic molecules • Homonuclear diatomic molecules – A more refined treatment of the MO theory considers interactions between the 2s and 2p orbitals – Orbital mixing causes the σs and σp molecular orbitals to move further apart in energy – The amount of mixing depends on the difference in energy between orbitals

5.7 Molecular orbital theory: diatomic molecules

Summary • Covalent bonds are formed as a result of the sharing of electrons between nuclei in a way that balances attractive and repulsive forces • Unequal sharing of electrons gives a polar covalent bond • Electronegativities generally increase across a period and decrease down a group

Summary • Ionic compounds are formed between elements with very different electronegativities • Lattice energy is the total energy that must be supplied to break the crystal lattice into its gaseous ions • Lewis structures show the distribution of valence electrons within a molecule

Summary • Lewis structures may be drawn using a five-step procedure • VSEPR theory states that a molecule will adopt a shape in which electron pair repulsions are minimised • A molecule with an overall asymmetric distribution of electrons has a dipole moment

Summary • Bond lengths in molecules are dependent on the atomic radii of the bonded atoms • Valence bond theory describes bonding in molecules using localised bonds formed from orbital overlap of hybrid orbitals • MO theory considers all possible overlaps between atomic orbitals in terms of delocalised bonds