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 nonmetallic 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 1s 2 2s 2 2p 6 3s 2 3p 4 1s 2 2s 2 2p 6 = core electrons = [Ne] [Ne] 3s 2 3p 4 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: ns 2 np 6 Ionic Bonding Na: 1s 2 2s 2 2p 6 3s 1 Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 + 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 = nonbonding important for geometry. 1

Group 7 7 valence electrons need 1 covalent bond for octet Group 6 6 valence electrons need 2 bonds for octet e.g. H 2 O Group 5: e.g. NH 3 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 (CH 4 ) 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 2 CO 2 N 2 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. SO 2 usually one with lowest electronegativity (most electropositive) exceptions H 2 O, 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 (HClO 4 ) Oxygen usually terminal (central in H 2 O) 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. IF 4+ - positive on I Anion place negative charge on most electronegative atom. e.g. ClF 4- - negative on F e.g. PCl 3 Place one pair of electrons between each pair of bonded atoms to form a single bond. 2

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 H 2 CO e.g. Sulphur trioxide, SO 3 Octet guideline Make single bonds all terminal atoms bound and each atom has a complete octet structure complete e.g. NH 3 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. NH 4+ all terminal atoms bonded = structure complete. If not octet of central atom must be extended. Isoelectric Species: How are NO +, N 2, CO and CN - similar? Multiple charge spread over as many atoms as possible not placed on single atoms alone. e.g. SO 4 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 (NO 2+ ) 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 atoms. 3

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 O 3. 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 (NO 3- ) 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. ClO 2, NO, NO 2, O - 2 odd valence electrons octet not achieved NO 2 : 5 + 2(6) = 17 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, BF 3, BCl; borax, Na 2 B 4 O 5 (OH) 4.8H 2 O BF 3 Could form octet by making a B F double bond NO: 5 + 6 = 11 e - Compounds i.e. BF 3 fairly reactive B can accommodate 4th e - - pair. Dative covalent (coordinate) 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. PCl 5 have to expand valence shell to place 10 e-s around central phosphorus atom. Also AsF 6-, SF 4, ICl - 4 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 (PO 4 3- ) 4

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. NH 3 4 electron domains around central atom 3 bonding pairs and one non-bonding pair. Each multiple bond = a single electron domain O 3 : 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 o apart) 3 e - domains trigonal planar (120 o apart) 4 e - domains tetrahedrally (109.5 o apart) 5 e - domains trigonal bipyramidal (120 o and 90 o apart) 6 e - domains octahedral (90 o apart) Shapes of AX n molecules / 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. BeF 2 ; trigonal planar (3 bonds) e.g. BF 3 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. nonbonding 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 3 e.g. BF 3 / CO 3 2- e.g. O 3 5

Four electron domains: Tetrahedral electron geometry AX 4 e.g. CH 4 e.g. NH 3 e.g. Cl 2 F + Molecules with expanded valence shell: Five electron domains Trigonal bipyramidal electron geometry AX 5 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 o from equatorial domain Each equatorial domain 120 o from other equatorial domains and 90 o from axial domains. Repulsions between domains at 90 o from 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. PCl 5 : SF 4 : ClF 3 : XeF 2 : 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. SF 6 : IF 5 : ICl 4- : Bond Angles Effect of non-bonding electrons and multiple bonds. Electron pair geometry of NH 3 is tetrahedral expect bond angle to be 109.5 o experimentally 107.5 o. Electron pair geometry of H 2 O 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. NO 3 ion bond angles = 120 o is this what is expected? 6