MPharm Programme PHA111 PDF

WEEK 9 MPharm Programme PHA111 Molecular Shape and Bonding Dr. Stephanie Myers Senior Lecturer in Medicinal Chemistry...

WEEK 9 MPharm Programme PHA111 Molecular Shape and Bonding Dr. Stephanie Myers Senior Lecturer in Medicinal Chemistry Dale 1.21 [email protected] Telephone: 0191 5152760 Slide 1 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 10 Pharmaceutical Chemistry Molecules, bonding and shape Functional groups alkenes, alcohols, alkyl halides, ketones, carboxylic acid derivatives, amines, aromatics Canvas - FUNCTIONAL GROUPS HANDOUT Acids and bases carboxylic acids, sulfonic acids, sulphonamides, imides, b- diketones, thiols, phenols amines (primary, secondary, tertiary, aromatic, heterocyclic), imines, amidines, guanidines Stereochemistry Kinetics Thermodynamics Slide 2 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Functional Groups Identity of functional group(s) in drug determines, for example: – route of administration – absorption from G.I.T. (if oral) – first pass metabolism – Bioavailability – binding to receptor – time course of effect (pharmacokinetics) Slide 3 of 68 MPharm PHA111 Molecular Shape and Bonding  Learning Outcomes You will learn the ability to: Consider chemical structure: Describe the difference between ionic, covalent bonds Appreciate the shape of chemical structures, hybridisation of atoms Identify functional groups Polar/non-polar, ionisable/neutral Deduce properties (solubility, absorption, distribution, metabolism, excretion) Understand how binding to a receptor produces a therapeutic effect Rationalise clinical care on the basis of medicinal properties and their side effects Design new medicines for specific purposes Give advice at all levels 4 Slide 4 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 A Pharmaceutical Chemistry Example: Atorvastatin Atorvastatin, a synthetic statin (used for controlling cholesterol) launched January 1997, patent expired 2011 in terms of cost, the most expensive (by gross ingredient cost) prescribed drug in 2008 largest selling small molecule drug in world : ~$10billion (2010); ~$15 billion peak (2007/8) 5 Slide 5 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atorvastatin Atorvastatin is a competitive inhibitor of HMG-CoA reductase, the rate-determining enzyme located in hepatic tissue that produces mevalonate lowers the amount of cholesterol produced which in turn lowers the total amount of LDL (low density lipoprotein) cholesterol also reduces blood levels of triglycerides slightly increases levels of HDL-cholesterol Slide 6 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Organic Chemistry All organic compounds contain the element carbon – 4A element – shares four electrons – forms four strong covalent bonds (or …. ?) – bonds to other carbons to create chains and rings – not all carbon compounds are derived from living organisms – over 99% of 37 million known compounds contain carbon Slide 7 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 The Periodic Table Slide 8 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Common Elements in Carbon- Containing Compounds Elements commonly found in organic compounds in the colours typically used to represent them Slide 9 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure Nucleus ‒ Positively charged ‒ Protons (+) and neutrons ‒ Small and dense (10-14 to 10-15 m in diameter) ‒ Contains essentially all the mass of an atom Electron Cloud ‒ Negatively charged electrons in a coud around the nucleus ‒ Atomic diameter is about 2Å ‒ 1Å = 10-10 m = 100 pm (picometers) Charge on proton = charge on electron ‒ Neutral atoms have equal numbers of protons and electrons ‒ If an atom gains an electron it becomes negatively charged ‒ If an atom loses an electron it becomes positively charged Slide 10 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic and Mass Numbers Atomic number, Z ‒ Number of protons in the atom's nucleus ‒ All atoms of a given element have the same atomic number Mass number, A ‒ Number of protons plus neutrons in the atom's nucleus ‒ Number of neutrons can be determined using M-Z Isotopes ‒ Atoms of the same element (same Z) that have different numbers of neutrons ‒ Therefore have different mass numbers (A) Mass number 12 13 14 6C 6C 6C Isotopes of carbon Atomic number 98.89% 1.11% trace magnetic spin = ½ radioactive 13C NMR t1/2 = 5730 years carbon dating Slide 11 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic and Molecular Mass Atomic mass (atomic weight) Ar ‒ Weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes Molecular Mass (molecular weight) Mr ‒ Sum of all atomic masses of all atoms in a molecule ‒ E.g. Ethanol has the formula CH3CH2OH (C2H6O) ‒ Mr = (2 x 12.011) + (6 x 1.008) + (1 x 15.9994) = 46.0694 atomic mass units (amu) Slide 12 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Calculating Moles A new definition of the mole was adopted by IUPAC in Jan 2018. The mole, symbol mol, is the SI unit of amount of substance. One mole contains exactly 6.02214076 × 1023 elementary entities. This number is the fixed numerical value of the Avogadro constant, NA, when expressed in mol−1, and is called the Avogadro number. Number of moles = mass of substance (g) / Molecular mass (Mr) E.g. Moles in 5.436g phenol (C6H5OH) = 5.436 g / 94.113 g/mol = 0.0578 mol Slide 13 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Periodic Table Trends Slide 14 of 68 MPharm PHA111 Molecular Shape and Bonding  Electronegativity Measures an atom's strength to attract electrons As you move to the right across a period of elements, electronegativity increases When the valence shell of an atom is less than half full, it requires less energy to lose an electron than gain one and thus, it is easier to lose an electron Conversely, when the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one As you move down a group, electronegativity decreases. This is because there is an increased distance between the valence electrons and nucleus The most electronegative element is fluorine (4.0) Slide 15 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Electronegativity Pauling Scale Graphic taken from Chemistry, Moore/Stanitski/Jurs, Thomson, 2006, pg 355 Slide 16 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals Electrons are not discrete particles Quantum mechanical model Can be described as both particles or waves Impossible to locate precise position of an electron Possible to indicate a region or volume where the electron is most likely to be found (1926, Erwin Schrödinger) ‒ This region is called an orbital ‒ Orbital denoted by Greek letter psi, Ψ ‒ Electron cloud has no sharp boundary Slide 17 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals Periodic table indicates which orbitals are present for each atom – in organic chemistry we are mostly concerned with s and p orbitals 1s 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 7p 4f 5f Slide 18 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals Four different kinds of orbitals for electrons: s, p, d, and f – s and p orbitals most important in pharmaceutical chemistry s orbitals - spherical, nucleus at centre p orbitals - dumbbell-shaped, nucleus at middle d orbitals - four cloverleaf-shaped and one dumbbell-doughnut (see slide 22) Slide 19 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals Orbitals are occupied by zero, one, or two electrons and are grouped in electron shells of increasing size and energy Electron Shell a group of an atom’s electrons with the same principal quantum number (same relative energy) – first shell; one s orbital, 1s, holds only two electrons – second shell; four orbitals, one s orbital (2s) + three p orbitals (2p), hold a total of eight electrons – third shell; nine orbitals, one s orbital (3s) + three p orbitals (3p) + five d orbitals (3d), hold a total of 18 electrons Slide 20 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals p Orbitals From the second shell onwards, there are three degenerate (same energy) mutually perpendicular p orbitals, px, py, and pz, of equal energy – they have symmetry about the x-, y- and z-axes, respectively Lobes of p orbitals are separated by region of zero electron density, a node – a node is a consequence of the wave-like properties – the node of a p atomic orbital is a plane passing through the nucleus, nodal plane zero probability of finding an electron at the p orbital nodal plane Slide 21 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Orbitals Adapted from :http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_At oms/Electronic_Orbitals Slide 22 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Atomic Structure: Electron Configuration Ground-state electron configuration Most stable, lowest-energy electron configuration of an atom Three rules: Aufbau principle lowest-energy orbitals fill first: 1s → 2s → 2p → 3s → 3p → 4s → 3d Pauli Exclusion Principle electron spin can have only two orientations (represented as up  and down ) only two electrons can occupy an orbital, and they must be of opposite spin to have unique wave equations Hund's rule if two or more empty orbitals of equal energy are available, electrons occupy each orbital with parallel spins until all orbitals have one electron Slide 23 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK Atomic number Element Electronic configuration 9 1 H 1s1 2 He 1s2 3 Li 1s2 2s 4 Be 1s2 2s2 5 B 1s2 2s2 2p1 6 C 1s2 2s2 2p2 1s2 2s2 2px1 2py1 7 N 1s2 2s2 2p3 1s2 2s2 2px1 2py1 2pz1 8 O 1s2 2s2 2p4 1s2 2s2 2px2 2py1 2pz1 9 F 1s2 2s2 2p5 10 Ne 1s2 2s2 2p6 11 Na 1s2 2s2 2p6 3s1 Energy 2 px py pz 2 px py pz 2s 2s 1s 1s Electronic configuration of Nitrogen Electronic configuration of Fluorine Slide 24 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Time for a Break Slide 25 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Orbital Hybridisation Carbon forms four bonds with other atoms – we will look at this in more detail later…. There is a mismatch in energy between the s and p electronic levels 2px 2py 2pz Energy 2s 4 sp3 orbitals 1s 1s Equivalent bonds require equivalent energy levels Energy is required to promote the electron s and p orbitals combine or ‘hybridise’ Produces four orbitals which have the same energy levels Slide 26 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 sp3 Hybridisation Video https://www.youtube.com/watch?v=RM3lO1xIAyE Slide 27 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Chemical Bonding Theory Atoms bond because the compound that results is more stable and lower in energy than the separate atoms ‒ Energy is released from the chemical system when a bond forms ‒ Energy is consumed by the system when a bond breaks Valence shell Outer most electron shell of an atom Eight electrons in valence shell (an electron octet) impart special stability to noble-gas elements in 8A Ne (2 + 8); Ar (2 + 8 + 8); Kr (2 + 8 + 18 + 8) Chemistry of main group elements governed by their tendency to take on electron configuration of the nearest noble gas Slide 28 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Chemical Bonding Theory Ionic Compounds Some elements achieve an octet configuration by gaining or losing electrons Ions form when an electron is gained or lost from a neutral atom Ions are charged because they have different numbers of protons and electrons Ions are held together by an electrostatic attraction, like in Na+ Cl-, forming an ionic bond Covalent Compounds Covalent Bond Bond formed by sharing electrons between atoms Molecule Neutral collection of atoms held together by covalent bonds Carbon achieves an octet configuration by sharing electrons Slide 29 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Ionic Bond Sodium (Na) has one 3s electron Loss of this electron leads to filled outer shell and formation of Na+ Energy is required to remove this electron - ionisation energy Na has a relatively low ionisation energy Elements with low ionisation energies are electropositive Chlorine (Cl) has seven valence electrons Gain of electron leads to filled outer shell and formation of Cl- Energy is released Elements that readily acquire electrons are electronegative The ions are independent species held together by electrostatic attraction Slide 30 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Covalent Bonds Covalent bond result from the overlap of two atomic orbitals Electrons are paired in overlapping orbitals and are attracted to nuclei of both atoms, thus bonding the two atoms together Non-polar covalent bonds are formed between atoms with similar electronegativity ( 1.7 Covalent bonds are formed when the electronegativity difference between the atoms is < 1.7 Non-polar covalent green > yellow > orange > red most positive most negative electrostatic potential electrostatic potential Slide 35 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Inductive Effect/Dipole Moment Inductive Effect An atom’s ability to polarise a bond Electron-withdrawing effect transmitted through s bonds Electronegative elements have an electron-withdrawing inductive effect Dipole Moment () A measure of the net polarity of a molecule Molecules as a whole are often polar Molecular polarity results from the vector summation (geometry) of all individual bond polarities and lone-pair contributions in the molecule Strongly polar substances are soluble in polar solvents, like water Slide 36 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Bond Representations Lewis structures (electron-dot structures) Dot representations of covalent bonds in molecules Valence shell electrons of an atom are represented as dots Kekulé structures (line-bond structures) Two-electron covalent bond is represented by a line Slide 37 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Representing Structures Slide 38 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Skeletal Structure / Line-Bond Structure A shorthand way of writing structures carbon atoms are assumed to be at each intersection of two lines (bonds) and at the end of each line Carbon atoms not usually shown, they are assumed Hydrogen atoms bonded to carbon are assumed Atoms other than carbon and hydrogen are shown The end of a line represents a carbon atom with 3 hydrogens, CH3 A two-way intersection is a carbon atom with 2 hydrogens, CH2 A three-way intersection is a carbon atom with 1 hydrogen, CH A four-way intersection is a carbon atom with no attached hydrogens 39 Slide 39 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Valency Valence electrons are those which occupy the outermost electron shell The no. of covalent bonds an atom forms depends on how many additional electrons are needed to reach a noble-gas configuration H (1s) needs one more electron to attain (1s2) and forms one bond C(2s22p2); four electrons to achieve filled outer shell (2s22p6); four bonds N (2s22p3); three electrons to achieve filled outer shell (2s22p6); three bonds Cl (2s22p5); one electrons to achieve filled outer shell (2s22p6); one bond The number of bonds an atom can combine with is referred to as a the valency of that element Slide 40 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Lone pairs Lone pair electrons Valence-shell electron pairs not used for bonding Also referred to as non-bonding electrons Filled p orbital Lone-pair electrons project out into space away from positively charged nuclei (directional based on hybridisation) giving rise to a considerable charge separation and contributing to the dipole moment Slide 41 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Assignment of Formal Charge Formal charge = number of valence e-s – (number of lone pair e-s + ½ number of bonding e-s) Slide 42 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Assignment of Formal Charge Carbon has four valence electrons: Formal charge = number of valence e-s – (number of lone pair e-s + ½ number of bonding e-s) Species where one atom has a single unpaired electron = a (free) radical Slide 43 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Assignment of Formal Charge The equation method to determine formal charge works well when simply looking at an isolated chemical structure, but when we begin to look at organic chemistry mechanisms it is quite impractical (time consuming!) Once you start drawing mechanisms – bond-making and bond-breaking processes - try to start thinking about the movement of electrons i.e. a neutral atom which gains two electrons will become negatively charged i.e. a neutral atom which loses two electrons will become positively charged i.e. a neutral atom which gains or loses one electron will become a radical Slide 44 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 Summary of Formal Charges Slide 45 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 9 To be continued…. Slide 46 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Valence Bond Theory Bonding theory that describes a covalent bond resulting from the overlap of two atomic orbitals Electrons are paired in overlapping orbitals and are attracted to nuclei of both atoms, thus bonding the two atoms together H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals H-H bond is cylindrically symmetrical Bonds formed by head-on overlap of two atomic orbitals along a line drawn between the nuclei are sigma (s) bonds Slide 47 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp3 Orbitals E p p p hybridisation sp3 sp3 sp3 sp3 s Combining one s and 3p orbitals Four equal orbital with ¼ s and ¾ p character Positive p lobe adds to s orbital generating larger lobe of sp3 hybrid Negative p lobe subtracts from s orbital generating smaller lobe of sp3 hybrid Larger lobe can overlap more efficiently with an orbital from another atom Formation of much stronger bonds Slide 48 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp3 Orbitals Slide 49 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp3 Bonding – Structure of Methane Slide 50 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp3 Bonding – Structure of Ethane Ethane C2H6 Two carbon atoms – both are tetrahedral C-H bonds formed by overlap of carbon sp3 and hydrogen s orbitals C-C bond formed by overlap of two carbon sp3 orbitals Increased bond length for C-C (versus C-H) Bond angles are near 109.5° σ bond Slide 51 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp2 Orbitals Combining one s and two p orbitals One electron is promoted from 2s to 2p Orbitals hybridise to 2sp2 3 equal orbitals (1/3 s and 2/3 p) and one unhybridised p orbital Slide 52 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp2 Orbitals Carbon can hybridise in other ways Two p orbitals can combine with the s orbital to give three sp2 orbitals One unhybridised p orbital remains Electron repulsion minimised by the three sp2 orbitals lying in a plane (flat) Bond angles all close to 120° – trigonal planar arrangement Unhybridised p orbital is perpendicular to the plane Slide 53 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp2 Bonding – Structure of Ethene The carbon atoms form two non-identical bonds with each other - a double bond Slide 54 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp2 Bonding – Structure of Ethene Double bond one bond formed by overlap of one sp2 orbital from each carbon (s bond) second bond formed by side-to-side overlap of the unhybridised p orbitals (p bond) four electrons are involved in the formation of a C=C bond C-C s bond is stronger than C-C p bond a C=C bond is stronger (730 kJ mol-1) and shorter than a C-C bond Large energy barrier to rotation (264 kJ mol-1) around the double bond (= strength of a p bond) This rotational barrier exists as the p orbitals must be well aligned for maximum overlap and formation of the p bond Slide 55 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp Orbitals Combining one s and one p orbital One electron is promoted from 2s to 2p Orbitals hybridise to 2sp 2 equal orbitals (1/2 s and 1/2 p) and two unhybridised p orbitals Slide 56 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp Orbitals One s and one p orbital combine to form two sp orbitals Two unhybridised p orbitals remain, perpendicular to sp orbital Bond angles 180o Linear arrangement Slide 57 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp Bonding – Structure of Ethyne Triple bond One bond formed by overlap of one sp orbital from each carbon (s bond) Second and third bonds formed by side-to-side overlap of the unhybridised p orbitals (p bond) Six electrons are involved in the formation of a C≡C bond Slide 58 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 sp Bonding – Structure of Ethyne Unhybridised p orbitals are perpendicular to each other and to the sp orbitals Two side-to-side overlaps give two π-bonds Region of high electron density around the C-C axis C≡C bond is stronger and shorter than a C=C bond Slide 59 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Bond Strengths, Bond Lengths and Bond Angles Bond Bond Bond Angle Molecule Bond Strength Length Shape (o) (kJ mol-1) (pm) methane, sp3 C-H 439 109 109.5 Tetrahedral CH4 ethane, sp3 C-C 377 154 109.5 Tetrahedral CH3CH3 sp3 C-H 420 109 ethene, sp2 C=C 728 133 Trigonal 120 H2C=CH2 sp2 C-H 464 108 Planar ethyne, HC sp C C 965 120 180 Linear CH sp C-H 558 106 Slide 60 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Bonding in the Methyl Cation Carbon is positively charged and bonded to three hydrogen atoms Three orbitals are hybridised – sp2 hybridisation The positively charged C and the three H atoms lie in a plane Non-hybridised p orbital remains empty and is perpendicular to the plane Slide 61 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Bonding in the Methyl Radical Carbon is bonded to three hydrogen atoms – carbon is sp2 hybridised The C and three H atoms lie in a plane – c.f. methyl cation Radical has one more electron than the cation – resides in the non-hybridised p orbital Slide 62 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation of Nitrogen Experiment shows bond angle in ammonia to be 107.3o – close to tetrahedral geometry – indication that nitrogen uses hybrid orbitals nitrogen hybridises to form four sp3 orbitals one of the four sp3 orbitals is occupied by two non- bonding electrons N-H bonds formed by sp3 – s overlap Slide 63 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation of Oxygen O-H bond formed by sp3-s orbital overlap (water, alcohols) C-O bond formed by sp3-sp3 orbital overlap (alcohols, ethers) Lone pairs occupy the remaining two oxygen sp3 orbitals Bond angle smaller than in methane (109.5o) Slide 64 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation of Nitrogen and Oxygen Hybridisation in multiply bonded nitrogen and oxygen is analogous to that of carbon: Slide 65 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation of Sulfur Commonly encountered in biological molecules – thiols … have a sulfur atom bonded to one hydrogen and one carbon – sulfides … have a sulfur atom bonded to two carbons Methanethiol, CH3SH – produced by some bacteria – simplest example of a thiol – sp3 hybridisation Dimethyl Sulfide (CH3)2S – simplest example of a sulfide – sp3 hybridisation Slide 66 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation of Phosphorus Most commonly encountered in biological molecules in organophosphates – compounds that contain a phosphorus atom bonded to four oxygens with one of the oxygens also bonded to carbon – ADP, adenosine diphosphate – ATP, adenosine triphosphate Methyl phosphate CH3OPO32- sp3 hybrid orbitals on phosphorus – tetrahedral geometry Slide 67 of 68 MPharm PHA111 Molecular Shape and Bonding WEEK 10 Hybridisation Summary ALL single bonds are s bonds ALL double bonds comprise one s bond and one p bond ALL triple bonds comprise one s bond and two p bonds To determine hybridisation state of a C, N, O, S atom, examine the number of p bonds it forms zero p bonds = sp3 hybridised one p bond = sp2 hybridised two p bonds = sp hybridised Exceptions: carbocations and carbon radicals – sp2 hybridised Phosphorus and sulfur are typically tetrahedral in shape and described as sp3 A p bond is weaker than a s bond – important! The greater the electron density in the region of orbital overlap, the stronger is the bond The more s character, the shorter and stronger is the bond The more s character, the larger is the bond angle Slide 68 of 68 MPharm PHA111 Molecular Shape and Bonding