BIOS4006 Topic 1 Wk3 2024-25 Moodle (2) PDF

Introduction to Biochemistry A Topic 1 Week 3 Chemical bonds - how atoms and molecules interact Type of bonds = ionic & covalent bonds = molecular formula for organic molecules = bond length & strength Lewis structur...

Introduction to Biochemistry A Topic 1 Week 3 Chemical bonds - how atoms and molecules interact Type of bonds = ionic & covalent bonds = molecular formula for organic molecules = bond length & strength Lewis structures to show paired and unpaired electrons = electron location & charge = bond polarity & resonance = bond angles and molecular shape Intermolecular forces = attractive forces between molecules Crowe & Bradshaw An extra recording will be available after your practical Chemistry for the Biosciences (4th Ed) = completing the Lab Notebook = data analysis and conclusions Parts of Chapters 2, 3, 4 & 8 = peer assessment activity and requirements From last week – we learnt about valency shell & electron numbers Valency 1 2 3 4 3 2 Outer Maximum 1 0 Electron valence shell electrons Shell 1 2 Shell 2 8 Shell 3 8 Shell 4 18 Elements in each GROUP (column) have the same number of valence electrons = they have the same number of electrons available to form bonds with other atoms = elements in each GROUP tend to behave similarly in chemical reactions From last week – we learnt about valency shell & electron numbers Valency 1 2 3 4 3 2 Outer Maximum 1 0 Electron valence shell electrons Shell 1 2 Shell 2 8 Shell 3 8 Shell 4 18 Elements in Groups 1 to 13 lose valence electrons more easily than gaining them Elements in Groups 15 to 17 gain electrons more easily than losing them Elements in Group 14 are Elements in Group 18 more likely to share electrons have full electron shells & than gain or lose them are unreactive  From last week – we learnt that electronegativity influences valency ELECTRONEGATIVITY = The tendency for an atom to attract the shared electrons in a chemical bond, when that atom is bound to others within a molecule More electronegative atoms attract electrons more strongly than less electronegative atoms - in a reaction, the electrons tend to move towards the more electronegative atom or group Caesium = least electronegative atom (current data) Fluorine = most electronegative atom I have researched the latest information since last week …… ELECTRONEGATIVITY = The tendency for an atom to attract the shared electrons in a chemical bond, when that atom is bound to others within a molecule More electronegative atoms attract electrons more strongly than less electronegative atoms - in a reaction, the electrons tend to move towards the more electronegative atom or group Francium = least electronegative atom (current Fluorine = most data) electronegative atom Chemical bonds = sharing of electrons (between atoms = intramolecular bond) + + + + + + + + Non-polar covalent Polar covalent Ionic bond Clean ionic bond bond bond Equality of electron share Difference in 0 1 2 3 electronegativity (between the atoms) Charge separation across the bond  Ionic bonds = electrostatic attraction between oppositely charged ions Formed by the (almost complete**) transfer of valence electrons between atoms with different electronegativity values (differences of 2.1 and above) Valence electrons (almost fully**) transfer from a weakly electronegative atom to a more electronegative atom The nuclei of both atoms attract the electrons The anion exerts the greatest pull on the valence electrons and is electrostatically attracted to the cation The opposite ionic charges leads to strong attraction between the ions = ionic bond The molecule overall has a neutral charge Na+ Cl – (ions are balanced) Consider the transfer complete for our purposes Characteristics of ionic bonds and the compounds they form Many are formed between atoms of metals and non-metals (large electronegativity difference ≥1.9) Good conductors of electricity - when molten, or in solution - not as solids High melting points (very stable) - ionic bonds are the strongest of all chemical bonds – require high energy input to be broken Good solubility in water - and other polar solvents Ionic bonds are POLAR – electrons are more attracted to one atom + + + + + + + + Non-polar covalent bond Polar covalent bond Ionic bond Clean ionic bond The overall molecule is neutral, but the two ions (dipoles) of the molecule have formal charges Direction of electron pull anion cation (increasing electronegativity) - + We say the molecule has Cl Na Dipole moment Negative Positive dipole dipole The polarity of ionic bonds causes attraction between molecules = intermolecular force The charged dipole of one molecule attract the dipoles of other molecules = ionic forces This is why many metal-nonmetal ionic compounds (salts) have crystalline structures Arrangement of ions in a lattice – rather than as discrete molecules - + - + + - + - Covalent bonds – valence electron sharing between atoms Formed by the sharing of valence electrons between atoms with similar electronegativity values (differences of less than 2.0) A covalent bond is the stabilising of attraction and repulsion forces between atoms The nuclei of both atoms attract the electrons The electrons minimise the repulsion of the nuclei and form a cloud that orbits a space around both atoms The molecule typically has a neutral charge (protons and electrons are balanced) Sharing 2 electrons forms a single covalent bond Sharing 4 electrons forms a double covalent bond Sharing 6 electrons forms a triple covalent bond Characteristics of covalent bonds and the molecules they form Form between non-metallic elements of similar electronegativity Poor conductors of electricity Variable melting points - Include molecules that are solids, liquids or gases at room temperature – require variable energy input to be broken variable solubility in water - And other solvents Covalent bonds can be polar or non-polar + + + + + + + + Non-polar covalent bond Polar covalent bond Ionic bond Clean ionic bond Non-polar bonds give rise to molecules that are neutral, with no charged dipoles (no dipole moment) Examples: O N O N Diatomic gases H2 N2 O2 Cl2 Hydrocarbons methane CH4 tetrachloromethane CCl4 Covalent bonds can be polar or non-polar + + + + + + + + Non-polar covalent bond Polar covalent bond Ionic bond Clean ionic bond Polar bonds also form neutral, but they have charged dipoles (dipole moment) due to uneven electron density The atoms linked by a polar covalent bond have a partial charge, indicated with the Greek letter delta (d) Examples: Chloromethane CH3Cl Water H20 Water is an important and unusual polar covalent molecule … Oxygen is more electronegative than hydrogen Water molecules therefore possess two single polar covalent bonds.. … and each atom has a partial charge, causing attraction between water molecules (intermolecular forces) More on this later ….. Water is an important and unusual polar covalent molecule … Oxygen is more electronegative than hydrogen Water molecules can break one of their single polar covalent bonds = dissociation or ionization … yields an hydrogen cation and a hydroxyl anion (hydroxide)  The ionization (dissociation) of water is an important concept to learn The dissociated hydrogen (H+) ion combines with another H20 molecule to form a Hydronium ion H30 + Water ionization therefore involves two molecules and there is no free hydrogen (H+) ion Water ionization occurs at a very low frequency, but constantly and reversibly – water is a mixture of H2O molecules and small amounts of the ions OH- and H3O+ – changes in the concentration of the H3O+ ion affects pH – changes in the concentration of the OH- ion affects pOH Some polyatomic ions are formed by covalent bonds Several of these “molecular” ions are important in biological processes and structures – worth learning several of them for this module and your general studies Cations Anions Ammonium NH4+ Hydronium OH3+ (H3O+) Hydroxide ion OH- Phosphate ion PO43- Nitrite ion NO2- Sulfate ion SO42- Carbonate ion CO32- Nitrate ion NO3- Calcium carbonate Ca CO3 Bicarbonate ion HCO3- Sulfite ion SO32- Ca2+ CO32- Found in limestone, chalks …… causes limescale and clay (Alps, Southern England) Covalent bond lengths and implications for bond strength Bond length is defined as the average distance between the nuclei of two covalently bonded atoms = the sum of the atomic radii of the two covalently bonded atoms Atomic radius = half the distance between the nuclei of two identical atoms, joined by a single covalent bond The bond length or atomic diameter of single-element dimers can be measured by x-ray diffraction atomic radius = half this atomic diameter Covalent bond length prediction Bond length can be predicted (calculated) from the atomic diameter of the relevant single element dimers If the atomic diameter (2R) of F-F = 142 pm and the atomic diameter (2r) of H-H = 74 pm What is the bond length (R+r) in molecule of H-F molecule atomic diameter atomic radius F-F 142 pm 71 pm H-H 74 pm 37 pm molecule predicted atomic diameter or bond length H-F (71 + 37 pm) = 108 pm Measured bond length for H-F = 92 pm Why is it different to the calculation..? Bond length depends on what atoms are bound together, and what type of covalent bond exists between them (single, double or triple) Covalent bond length is inversely related to bond strength Bond dissociation energy or ionization energy is the energy required to break a covalent bond. In longer bonds, the shared electron pair is further from the atomic nuclei In general, longer bonds require lower bond dissociation energy to be broken, than shorter bonds Bond Bond Bond Bond length energy length energy (pm) (kJ/mol) (pm) (kJ/mol) H-H 74 436 O-O 148 145 C-C 154 348 O=O 121 498 C=C 134 614 F-F 142 158 C=C 120 839 Cl-Cl 199 243 N-N 145 170 Br-Br 228 193 N=N 110 945 I-I 267 151 Try some of these for yourself over the break Bond length e Calculated (pm) (k bond length H-H 74 (pm) C-C 154 H-C C=C 134 H-N C=C 120 Bond H-O N-N 145 length e N=N 110 (pm) (k H-F O-O 148 H-Cl O=O 121 H-Br F-F 142 H-I Cl-Cl 199 C-N Br-Br 228 C-O I-I 267 C-F C-Cl C-Br C-I There are conventions for representing for covalent bonds (organic molecules) Empirical formula – ratio of atoms in the molecule, determined experimentally CnH2n Gives no indication of actual numbers of atoms in the molecule (not useful in biology) CnH2n +2 Molecular formula – correct numbers for each atom (or group) in the molecule May give some indication of the arrangement of the atoms in the molecule In organic molecules this might be very different C2H6O = – there are 5 possible isomers of this formula C2H5OH hexane 2-methyl pentane C6H14 3-methyl pentane 2,2-dimethyl butane 2,3-dimethyl butane Structural formula is useful to reveal covalent bond positions Structural formula gives the correct (full) molecular formula, with all covalent bonds illustrated Single covalent bonds are drawn as a single line Double and triple covalent bonds are drawn as double and triple lines Ethane C2H7 Carbon dioxide CO2 Hydrogen cyanide HCN Ethanol C2H5OH OH Structural formula is essential to reveal isomers This approach distinguishes the 5 isomers of the molecular formula C6H14 hexane 2,2-dimethyl butane 2,3-dimethyl butane 2-methyl pentane 3-methyl pentane Condensed structural formula hides C-H bonds Condensed (abbreviated) structural formula – correct (full) molecular formula, but (a) all bonds between C and H atoms must be assumed (b) only shows covalent bonds between carbon and non-hydrogen atoms (c) can be drawn without bond lines, using brackets to indicate where branched groups bond hexane CH3 CH2 CH2 CH2 CH2 CH3 CH3CH2CH2CH2CH2CH3 CH3 CH3 CH CH2 CH2 CH3 2-methyl pentane CH3CH(CH3)CH2CH2CH3 CH3 3-methyl pentane CH3 CH2 CH CH2 CH3 CH3CH2CH(CH3)CH2CH3 Condensed structural formula hides C-H bonds Condensed (abbreviated) structural formula – correct (full) molecular formula, but (a) all bonds between C and H atoms must be assumed (b) only shows covalent bonds between carbon and non-hydrogen atoms (c) can be drawn without bond lines, using brackets to indicate where branched groups bond CH3 2,2-dimethyl butane CH3 C CH2 CH3 CH3C(CH3)2CH2CH3 CH3 CH3 2,3-dimethyl butane CH3 CH CH CH3 CH3CH(CH3)CH(CH3)CH3 CH3 Hydrogen cyanide CH=N = Skeletal formula offers minimal representation, hiding C and H atoms Skeletal formula – minimal representation showing only covalent bonds (a) Each line represents a single bond between 2 carbon atoms (b) Other atoms (oxygen, nitrogen) would be identified by their element symbol How do we know how many “hidden” hydrogens attach to each carbon (or other atom) …? Hint – it’s linked to valency (c) Double and triple covalent bonds would be indicated with a double or triple line ==N == Gilbert Lewis (1916) proposed the OCTET rule to predict (most) covalent bonding Atoms seek to share enough electrons to complete their valence shells The most important covalent bonds in organic molecules are between elements with valence electrons in shells 2 and 3 = maximum capacity of 8 (Hence the OCTET rule) The rule doesn’t work for heavier elements Gilbert Lewis – proposed the OCTET rule to predict covalent bonding Covalent bonds are formed by the sharing of a pair of electrons Electron pairs involved in forming covalent bonds are termed bonding pairs A single covalent bond consists of one bonding pair of electrons (a double bond consists of 2 bonding pairs, etc) The valency of each element predicts how many bonding pairs will fill their valence shell, and hence how many covalent bonds an atom can form Ethene C2H4 Methane CH4 Ethane C2H6 bonding pairs Hydrogen = 1 valence electron in shell 1 Carbon = 4 valence electrons in shell 2 Needs 2 electrons to fill shell 1 Needs 8 electrons to fill shell 2 = will share 1 bonding pair of electrons = will share 4 bonding pairs of electrons = will form 1 covalent bond = will form 4 covalent bonds Some valence electron pairs are only rarely shared as covalent bonds Electron pairs of more electronegative atoms are not always shared with another atom Electron pairs that are NOT involved in forming covalent bonds are termed non-bonding pairs or lone pairs Hydronium H3O+ Water H2O lone pairs H bonding pairs Oxygen = 6 valence electrons in shell 2 Needs 8 electrons to fill shell 2 = can share 4 bonding pairs of electrons = can form 4 covalent bonds BUT oxygen normally forms 2 covalent bonds - 2 electron pairs are lone pairs 1 lone pair can be used to form another covalent bond with a cation Some valence electron pairs are only rarely shared as covalent bonds Electron pairs of more electronegative atoms are not always shared with another atom Electron pairs that are NOT involved in forming covalent bonds are termed non-bonding pairs or lone pairs Ammonia NH3 lone pair Ammonium NH4+ H bonding pairs Nitrogen = 5 valence electrons in shell 2 Needs 8 electrons to fill shell 2 = can share 4 bonding pairs of electrons = can form 4 covalent bonds BUT nitrogen normally forms 3 covalent bonds - 1 electron pair is a lone pairs The lone pair can be used to form another covalent bond with a cation Lewis Symbols help display valence electron numbers within atoms Lewis symbols use strong dots to show the valence electrons of an atom (core electrons not shown) This helps us predict: - how many covalent bonds can form between atoms = bonding / shared electron pairs - any uneven distribution of electrons (dipole moment) - the shape of a molecule (bond angles) Electrons are shown at the 4 “compass points” Octet rule – maximum of 8 - shown in 4 pairs Lewis Structures help display electron pairs within covalent molecules Lewis symbols are combined to predict the Lewis structure for a covalent Electrons form covalent bonds where they are shared in pairs between atoms = The bonding electron pairs Logical structure is 4 single covalent bonds Carbon atom has 4 = 4 shared electron pairs valence electrons and valency = 4 + CHECK the structure leads 1 shared bonding electron Hydrogen atoms have 1 to full valence shells (Lewis’s pair = single covalent bond valence electron and valency = 1 octet rule) Non-bonding electron pairs are also shown in Lewis structures Valence electrons that do not form covalent bonds must also be shown in Lewis structures These are non-bonding electron pairs or lone pairs 7 valence electrons 1 shared electron pair 1 single covalent bond valency = 1 3 lone electron pairs 3 lone electron pairs on the bromine atom on the bromine atom 3 shared electron pairs 1 triple covalent bond 5 valence electrons 1 lone electron pair on 1 lone electron pair on valency = 3 each nitrogen atom each nitrogen atom Lone (non-bonding) electron pairs can be shared with ions to make another bond We’ve already seen an example of this: lone electron pair Ammonia NH3 + Ammonium NH4+ The lone electron pair of the nitrogen atom can be shared with a proton (H +) to form a covalent bond + + The result is a polyatomic ion, with a formal (overall) charge of 1+ Lone (non-bonding) electron pairs can be mobile in some molecular ions For polyatomic ions with lone electrons, we can propose different, equally correct Lewis structures known as resonance structures Example: carbonate ion, CO32- In each form of the Lewis structure for a carbonate ion: The central carbon atom has 4 bonds (= full valence shell) ONE oxygen atom has a double covalent bond + 2 lone pairs All the oxygen atoms have a full valence shell TWO oxygen atoms have a single covalent bond + 3 lone pairs = access to 8 electrons Both have a formal charge (there is 1 more electron than the number of protons) Lone (non-bonding) electron pairs can be mobile in some polyatomic ions Each resonance structure exists only fleetingly – the non-bonding electrons constantly move across the whole molecule – the electrons are delocalized The “average” representation is more important = the resonance hybrid = The resonance hybrid has a formal charge Each bond is a hybrid between a single and double covalent bond All three oxygen atoms possess a share of the overall negative charge Other examples of molecules with resonance hybrids Electrons are delocalized in the polar molecular phosphate ion PO43- = Electrons are also delocalized in the non-polar, cyclic (aromatic) molecule benzene C6H6 = = = Be aware of the limitations of Lewis structures We need to remember the relative electronegativities of the atoms sharing electrons Not all electron sharing is equal – polar covalent bonds have partial charges (but not formal charges) Molecules are 3-dimensional whereas chemical structures are 2-dimensional Molecules are not ‘static’ – atoms vibrate and parts of the molecules can ‘rotate’ with respect to other parts In some molecules, the bonding network framework is not fixed and some of the bonds between atoms are in a state of flux The octet rule is not always applied exactly Examples the phosphate ion (PO43- ) – the central phosphorus forms 5 covalent bonds - the atom therefore has 10 valence electrons - these are accommodated in the 3-d orbital Try these for yourself over the break Give the condensed (abbreviated) formula structure Draw the Lewis dot structure HCN CO2 Draw the skeletal structure CH3CH(CH3)CH2CH3 Are these Lewis structures correct? CH3CH(OH)CH2CH2CH3 The valence electrons influence bond angles and molecular structure Bond angles consider the relative spatial arrangement of the covalent bonds around the central atom Valence shell electron pair repulsion (VSEPR) model All electron pairs distribute themselves around the central atom to minimise electrostatic repulsion = maximize the distance between them CO2 BF3 CH4 The VSEPR model predicts bond angles and molecular structure Bond angles depend on the number of electron pairs surrounding the central atom Electron pairs CO2 BF3 CH4 PCl5 SF6 But why is the shape of water different to carbon dioxide? The central carbon atom of CO2 has 4 pairs of The central oxygen atom of H2O has 4 pairs of electrons electrons Only 2 are bonding pairs (2 single covalent bonds) All are bonding pairs (2 double covalent bonds) Oxygen typically has 2 non-bonding pairs Maximum repulsion = straight line (180◦ angle) These influence the structure as though they were 2 single covalent bonds Maximum repulsion = tetrahedral (like CH4) Lewis structure for CO2 Lewis structure for H2O (carbon typically has no lone pairs) This also explains the shape of ammonia – another tetrahedron The central nitrogen atom of NH3 has 4 pairs of electrons Only 3 are bonding pairs (3 single covalent bonds) Nitrogen typically has 1 non-bonding pair This again influences the structure as though they were a single covalent bond Maximum repulsion = tetrahedral (like CH4) Lewis structure for NH3 Molecules are attracted to each other by intermolecular forces Ionic and covalent bonds, which link atoms within a molecule, are intramolecular forces Much weaker – but usually lots of them and hence important = intermolecular forces exist between molecules We’ve seen one already in ionic compounds = ionic interactions Formally charged ions give the neutral ionic molecules two dipoles… … which attract the dipoles of other molecules - + - + + - + - Covalent molecules with formal charge also show ionic interactions Covalent molecules (polyatomic ions) with a formal charge also have dipoles that mediate attractive forces One of the most important examples of this in biology is the interaction between distant amino acids in a protein molecule A polar carbonyl group (CO2-) is attracted to a polar amino group (NH3+) The 3-D structure of proteins, fats and carbohydrates is a result of both intramolecular bonds and many / various intermolecular forces - you’ll learn more about these in the lectures on Topic 3 Covalent molecules with partial charges induce a different attraction Covalent molecules with polar bonds possess partial charges (dipoles) These can attract similar dipoles leading to intermolecular forces known as Dipole interactions Molecular dipoles can result in steric repulsion as well as attraction Steric repulsion acts when dipoles of the same charge are brought close together Temporary dipoles produce attraction known as dispersion forces Temporary dipoles are due to the influence of other atomic nuclei on the electron cloud of a molecule The nucleus of one atom can form an electrostatic attraction with the electrons of another atom This causes a temporary distortion of the electron cloud = a temporary dipole moment Hence there is a temporary intermolecular attraction = dispersion force or London force London / dispersion forces and dipole interactions are forms of Van der Waal forces Hydrogen bonds are a significant intermolecular force in biology Hydrogen bonds form where a hydrogen atom is covalently bonded to an electronegative atom … … giving the covalent molecule dipole moment (uneven electron distribution) Example – H2O The strongly electronegative oxygen atom forms the negative dipole of a water molecule The weakly electronegative hydrogen atom forms the positive dipole in water molecules Hydrogen and oxygen dipoles attract via non-bonding electron pairs The lone pairs (non-bonding pairs) of electrons on the oxygen atom will attract the positive dipole of another water molecule = Oxygen is the hydrogen bond donor The hydrogen will be attracted to the lone pairs (non-bonding pairs) of electrons in the oxygen atoms of other water molecules = Hydrogen is the hydrogen bond acceptor Hydrophobic interactions occur when polar molecules disrupt water Hydrophobic interactions Non-polar molecules (or regions of molecules) cannot form hydrogen bonds with water Rearrangement of water to minimise disruption of hydrogen bonding results in a “cage” of water around aggregations of the polar molecules You’ll need to read sections from different chapters of our textbook Section 2.7 Lewis symbols for atoms Chapter 3 Chemical bonds Especially: 3.2 electronegativity 3.3 ionic bonds 3.4 & 3.5 covalent bonds Toolkit 3 Lewis structures Section 8.1-8.2 bond angles and lengths Chapter 4 intermolecular forces Sections 4.1 - 4.4