Theme 3: Bonding and Molecular Structure. (Chapter 8)

Theme 3: Bonding and Molecular Structure.

(Chapter 8)

End of Chapter questions:

5, 7, 9, 12, 15, 18, 23, 27, 28, 32, 33, 39, 43, 46, 67, 77

Chemical reaction – valence electrons of atoms rearranged (lost, gained, reorganised) to form a net attractive force → chemical bond between atoms.

• Ionic bond – one or more valence electrons transferred from one atom to another (ions have noble gas electron configuration) → creating positive and negative ions, → electrostatic attraction between positive and negative ions.

• generally interaction of metals and non-metals.

• Covalent bonding – involves sharing of electrons between 2 atoms, i.e. interaction between non- metallic elements.

• Generally electrons are shared unequally – from very little (ionic) to considerable (covalent).

Lewis Symbols

• 2 types of electrons in atoms, i.e.

• core electrons that are not involved in bonding or chemical reactions, and

• valence electrons, i.e. electrons in the outermost electron shell of an atom, involved in bonding.

• Periodic Table Main groups: Group number = number of valence electrons, i.e. electrons in the s- and p-electrons in the highest n.

• Transition metals – valence electrons include ns + (n-1)d electrons.

• Remaining electrons = core electrons.

Lewis Dot Symbols

Useful way to represent valence electrons of atom.

• Element symbol – represents nucleus and core electrons.

• Valence electrons – represented by dots

• 1st 4 placed around symbol

• Additional electrons paired with those already there.

e.g. S: 16 electrons 1s²2s²2p⁶3s²3p⁴

• 1s²2s²2p⁶= core electrons = [Ne]

• [Ne] 3s²3p⁴ 6 valence electrons (group 6)

Main Group: valence shell can accommodate 4 pairs of electrons = 8 electrons →

→

• Nobel gases (except He) have 8 electrons – unreactive / inert.

• Octet Rule: Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence e^-s.

• Octet = full s- and p- subshells in an atom:

ns²np⁶

• Ionic Bonding Na: 1s²2s²2p⁶3s¹→ Cℓ: 1s²2s²2p⁶3s²3p⁵+ e^-

Covalent Bonding:

• bonds result when one or more e^-- pairs are shared between 2 atoms.

• i.e. H· + ·H → H:H each H-atom acquires 2nd electron → achieves stable, noble gas configuration of He.

• Cℓ - Cℓ : by sharing the bonding pair each Cℓ atom has 8 electrons (an octet) in valence shell.

• Lone pairs – not involved in bonding = non- bonding

• important for geometry.

• Group 7 → 7 valence electrons – need 1 covalent bond for octet

• Group 6 → 6 valence electrons – need 2 bonds for octet

• e.g. H

O

• Group 5: e.g. NH

Questions:

1. Predict the formula of a stable binary (2 elements) compound when nitrogen reacts with fluorine.

2. Compare the Lewis symbol for Neon with the Lewis structure of methane (CH₄) – how are they similar / different?

Multiple Bonds

• Sharing one pair of electrons – single covalent bond.

• Many molecules – share more than one pair of electrons to obtain an octet

between 2 atoms = double or triple bonds.

• e.g. O

• CO

• N

Bond Length

• General rule: bond length decreases as number of shared electrons pairs increases.

N – N N ═ N N ≡ N

1.47 Å 1.24 Å 1.10 Å

• Lone pairs – in same valence shell as bonding electrons, influence molecular geometry.

Drawing Lewis electron – dot Structures.

Determine arrangement of atoms within molecule.

• Central atom – for simple compounds, often 1st atom is central e.g. SO₂

– usually one with lowest electronegativity (most electropositive)

– exceptions H₂O, common acids (acidic H first)

• Hydrogen – only one electron, can only make one bond = terminal position.

• Halogens – often terminal atoms, forming one bond

– can be central when combined with oxygen in oxoacids (HCℓO₄)

• Oxygen usually terminal (central in H₂O)

2. Write symbol of each element

• indicate valence electrons of each atom using dots/crosses

• cation → place positive charge on most electropositive atom. e.g. IF₄⁺- positive on I

• Anion → place negative charge on most electronegative atom.

• e.g. ClF₄^-- negative on F

• e.g. PCℓ₃

• Place one pair of electrons between each pair of bonded atoms to form a single bond.

• Central atom – if less than 8 electrons, use lone pairs on terminal atoms as bonding pair for multiple bond. (Multiple bonds usually with C, N, O).

• e.g. Formaldehyde H₂CO

• e.g. Sulphur trioxide, SO3

Octet guideline …

• Make single bonds → all terminal atoms bound and each atom has a complete octet → structure complete e.g. NH₃

• If not → make double bond. Check for octet.

• If not → triple bond.

• Central atom has complete octet but terminal atoms not yet bonded → dative covalent bond from lone pair on central atom, e.g. NH₄⁺→ all terminal atoms bonded = structure complete.

• If not → octet of central atom must be extended.

• Isoelectric Species: How are NO⁺, N₂, CO and CN^-similar?

Multiple charge …

• → spread over as many atoms as possible – not placed on single atoms alone.

• e.g. SO₄^2-

• S – most electropositive element → central atom

• First make single bonds and account for lone pairs before determining whether multiple bonds are present.

• e.g. Nitronium ion (NO₂⁺)

Formal Charge

• Valence electrons are not evenly distributed, as suggested by Lewis structures. Electron pairs may be drawn more strongly towards one atom in a bond → making that atom slightly negative (δ-).

• The way electrons are distributed in a molecule, i.e. charge distribution, affects the properties of molecules. e.g. H – Cℓ

δ+ δ-

• δ+ end of one molecule is attracted to the δ- end of another molecule

→ intermolecular forces affect properties of substance, i.e. boiling pt

Formal charge = Group number of atom – [LPE + ½ (BE)]

• Group number = valence electrons of atom

• LPE = lone pair electrons

• BE = bonding electrons …electrons assigned by Lewis diagram

• Concept of formal charge helps to determine which Lewis structure is most valid – not real charge, a form of ‘bookkeeping’

• Atom → positive if ‘contributes’ more electrons than it ‘gets back’.

→ negative if contributes fewer electrons

→ lone pairs belong to the atom to which they are allocated

→ bonding electrons are shared between bonded atoms ( ½ BE)

→ sum of formal charges on atoms in molecule equals the pos/neg charge on molecule/ion

• e.g. OH^-

• Generally choose Lewis structure in which atoms bear formal charges closest to zero

• Negative charges on more electronegative

Resonance

• Resonance structures are used to represent bonding in a molecule/ion when a single structure fails to describe actual electronic structure accurately.

• e.g. ozone O₃.

• Ozone bonds – length 127.8 pm

• Usually average O – O 132 pm and O ═ O 121 pm

• Both O – O bonds are the same length → implies bonds are equivalent.

• Bond angle 116.8^o< 180^o(due to presence of lone pair on central atom)

• Double headed arrow (↔) indicated the true structure is between the 2 extremes.

• e.g. Nitrate ion (NO3-) – 3 equivalent Lewis structures.

• Arrangement of atoms is the same in each structure, only placement of electrons differs.

• Electronic structure is a hybrid of all 3 resonance structures.

• Formal charge on oxygen atoms = -²/₃

i.e. total of formal charges on O = -1 + -1 + 0 = -2, divided by the number of O-atoms = 3

Exceptions to the Octet Rule

1. Molecules / ions containing odd numbers of electrons:

• e.g. CℓO₂, NO, NO₂, O₂^- odd valence electrons – octet not achieved

• NO₂: 5 + 2(6) = 17 e^-

• NO: 5 + 6 = 11 e^-

2. Molecules / ions in which the central atom has fewer than 4 pairs (8) of electrons.

• e.g. Boron – 3 valence electrons, forms 3 bonds → valence shell with 6 electrons.

• Forms many compounds i.e. boric acid, B(OH)₃; boron trihalides, BF₃, BCℓ; borax, Na₂B₄O₅(OH)₄.8H₂O

• BF₃Could form octet by making a B – F double bond

• Compounds i.e. BF₃fairly reactive – B can accommodate 4th e^-- pair. Dative covalent (co- ordinate) bond can be formed.

3. Molecules / Ions where the central atom has more than an octet (8) of valence e-s

– Elements in 3rd and higher periods

• have ns and np and unfilled nd –orbitals that can be used for bonding.

• e.g. PCℓ₅ – have to ‘expand’ valence shell to place 10 e-s around central phosphorus atom.

• Also AsF₆^-, SF₄, ICℓ₄^-

• Elements of 2nd period are restricted to a maximum of 8 electrons in their compounds.

• Expanded valence shells occur most often when central atom is bonded to small and highly electronegative atoms, F, Cℓ, O.

• Some Lewis structures are written with an expanded valence shell, even though they can be written with an octet (giving a better group of formal charges).

• e.g. phosphate (PO₄^3-)

Molecular geometry.

• Single central atom bonded to 2 or more atoms (of the same type) AX_n → A = central atom

X = terminal atom/s

• Possible shapes depend on value of n.

• A – main group element (s- and p- block) use VSEPR (valence shell electron-pair repulsion) model → bonding and lone pair electrons in the valence shell of an element repel each other and seek to be as far apart as possible.

• Bonding pair e^-s= defines a region in which electrons are most likely to be found = electron domain

• Non-bonding pair e^-s(lone pairs) = electron domain – located principally on one atom.

• e.g. NH₃

• 4 electron domains around central atom – 3 bonding pairs and one non-bonding pair.

• Each multiple bond = a single electron domain

• O₃:

• Central atom: 3 e^-domains – 2 bonding and 1 non-bonding.

• Electron domains – negatively charged, therefore repel each other and try to be as far apart as possible.

• Best arrangement of electron domains is the one that minimises the repulsions between them:

• 2 e^-domains – arranged linearly (180^oapart)

• 3 e^-domains – trigonal planar (120^oapart)

• 4 e^-domains – tetrahedrally (109.5^oapart)

• 5 e^-domains – trigonal bipyramidal (120^o and 90^oapart)

• 6 e^-domains – octahedral (90^oapart)

• Shapes of AXnmolecules / ions depend on the number of e^-domains surrounding central A- atom.

• Electron pair geometry – geometry of all valence electron pairs around central atom.

• Molecular geometry = bonding geometry – arrangement in space of central atom and terminal atoms.

• Simplest VSEPR – electron pairs around central atom involved in single covalent bonds.

• Linear (2 bonds) e.g. BeF₂;

• trigonal planar (3 bonds) e.g. BF₃– central atom – no octet.

• Lone pairs on the central atom occupy spatial positions even though they are not included in the description of the shape of the molecule / ion.

• VSEPR model → predict electron domains geometry → if all electrons domains are bonded

→ molecular geometry = electron geometry.

• Draw Lewis structure of molecule / ion → count electron domains around central atom (i.e. non- bonding pairs, single bonds, double bonds)

• Determine electron domain geometry → arrange electron domains around central atom to minimise repulsions.

• Use arrangement of bonded atoms to determine molecular geometry.

• Trigonal Planar: AX

e.g. BF

/ CO

• e.g. O

• Four electron domains: Tetrahedral electron geometry AX

• e.g. CH

• e.g. NH

• e.g. Cℓ

F

Molecules with expanded valence shell:

• Five electron domains – Trigonal bipyramidal electron geometry AX₅

• 2 sets of positions that are not equivalent → 3 equatorial electron domains define and equatorial triangle

• → 2 axial electron domains – north and south

‘poles’ of molecule

• Each axial domain – 90^ofrom equatorial domain

• Each equatorial domain – 120^ofrom other equatorial domains and 90^ofrom axial domains.

• Repulsions between domains at 90^ofrom each other are greater than when angle is 120^o.

→ Equatorial domain experiences less repulsion

→ Lone pairs exert larger repulsions – will occupy equatorial positions

→ double bonds have greater electron density → more repulsion → equatorial position.

• If terminal atoms differ → more electronegative in axial positions (bonds will be longer).

• e.g. PCℓ₅:

•

• SF₄:

•

• CℓF₃:

•

• XeF₂:

Six electron domains – most stable geometry is octahedral (6 vertices)

• All bond angles are 90^o

• All six vertices are equivalent

→ non-bonding domain can be placed in any position

→ 2 non-bonding domains – placed opposite each other to minimise repulsion

• e.g. SF₆:

•

• IF₅:

• ICℓ₄^-:

Bond Angles – Effect of non-bonding electrons and multiple bonds.

• Electron pair geometry of NH₃is tetrahedral – expect bond angle to be 109.5^o→

experimentally 107.5^o.

• Electron pair geometry of H₂O also tetrahedral

→ experimental bond angle 104.5^o

• Bond angle decreases with increasing number of non-bonding electrons.

• Bonding electron pair attracted by both nuclei of bonded atoms.

• Lone pair attracted by only one nucleus → electron domain more spread out → exert a greater repulsive force on adjacent electron domains (compress bond angles).

Multiple bonds also contain a higher electron charge density than single bonds → also larger electron domains.

• Relative strength of repulsions: Lone pair – lone pair > lone pair – bonding pair > bonding pair – bonding pair

• Q. NO₃⁻ion bond angles = 120^o…is this what is expected?