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Section ALesson 1 – Structure of Atoms, Electrons and Orbitals Electronic Configuration = Number of electrons an atom has and how they are arranged around the nucleus Aufbau Principle = Lowest energy orbitals are occupied first Pauli Principle = Only 2 electrons occupy each orbital Electrons and Orb...

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Section A Les Lesson son 1 – Structure of Atoms, Electrons and Orbitals Electronic Configuration = Number of electrons an atom has and how they are arranged around the nucleus • Aufbau Principle = Lowest energy orbitals are occupied first • Pauli Principle = Only 2 electrons occupy each orbital Electr Electrons ons and Orbitals • Fundamental Property of Quantum Mechanics = Matter has wave-like properties (wave-particle duality) • Wave-Particle Duality = Electrons are not just a particle and not just a wave •

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Wavefunction = A mathematical description of the distribution of electrons in terms of position and time

The further the orbital is from the nucleus, the higher the energy Wavefunctions can be used to calculate the radial distribution function of the orbital (how the probability of finding the electron varies with distance from the nucleus) o For the 1s orbital the electron has the highest probability of being found at the dotted line o It has a 0 probability of being found at the nucleus, but the chances increase rapidly with r, until the maximum o The probability of the electron being found a long, long way from the nucleus is very low – tends to zero o r = probability of electron being found at that distance o The energy of the orbital depends on the strength of the nuclear charge that the electron feels when it occupies that orbital o The 1s orbital would shield the 2s orbitals therefore the probability of an electron being further away from the nucleus for the 2s orbitals is higher Uncertainty Principle = Position and momentum of electron cannot simultaneously be determined Orbital = The most probable position of an electron (region of space) o o

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The Four Quantum Numbers 1. Principle Quantu Quantum m Number (n) = Size of orbital or how far an electron occupying it is from the nucleus o Shells are of increasing size o Each electron in an atom is labeled with 4 quantum numbers 2. Angular Momentum Quantum Number (l) = Shape of orbital o l = 0, …, n-1 o l = 0 => s orbital o l = 1 => p orbital 3. Magnetic Quantum Number (ml) = Orientation of orbital

o ml = -l, …, +l o x, y, z orientations 4. Spi Spin n Quant Quantum um Number ((m ms) = Electron spin o Can have a principal axis and a rotation axis that can spin around o It can spin in 2 directions o ms = +1/2 or –1/2 o +1/2 = Clockwise rotation around the axis of the electron o –1/2 = Anticlockwise rotation around the axis of the electron •

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Pauli Exclusion Principle = No 2 electrons in an atom can have the same 4 quantum numbers o Implication = A maximum of 2 electrons can share each orbital, with the same values of n, l, ml but a different ms Hund’s Rule of Maximum Multiplicity = If 2 or more orbitals have the same energy, electrons will spread out to occupy the maximum possible number of orbitals, maximizing the number of parallel spins If electrons occupy separate orbitals, then they occupy different regions of space so there is less electrosta electrostatic tic repulsion – lower energy arrangement Spin Correlation = Electrons have lower energies if their spins are parallel because parallel spins will stay further away from each other so repulsion is reduced

Les Lesson son 2 – Covalent Bonding: Lewis and VSEP VSEPR R

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When 2 atoms move closer to each other, the potential energy of the atoms decreases Potential energy between 2 atoms varies with the distance between their nuclei Lewis Model = Covalent bonding occurs when valence electrons are shared between 2 atoms Octet Rule = Each atom acquires shares in electrons until its valence shell achieves 8 electrons

Lewis Model of Ions = Molecules can accept or lose electrons to give a complete valence shell Coordinate Bonds = Both electrons come from one atom to form a covalent bond Resonance Hybrids = Possible to draw more than one satisfactory Lewis structure for a particular molecule • Hypervalent Compounds = Compounds that require more than an octet of electrons in order to draw a Lewis structure Valence She Shellll Electr Electron on Pair Repulsion Theory (VSEPR Theory) • Basic Assumptions o Electrons in bonds and lone pairs around an atom can be considered as charge clouds which repel each other o The lowest energy arrangement is when the atoms are as far apart as possible, and this determines the equilibrium molecular shape o Lone pairs repel more than bonding pairs o If a lone pair has a choice between an equatorial position and an axial position, it will occu py the equatorial site This is because in the equatorial position it is repelled less by the 2 axial bonding pairs than it would be by the 3 equatorial bonding pairs if it was in the axial position • Importance of Geometry and Bonding o Active sites of enzymes specific for substrate of particular size and shape => greater selectivity o Cell receptors/signaling pathways/responses to hormones depend on molecular recognition o Antibody-antigen interactions o If the wrong molecule binds to an enzyme active site this can lead to problems such as inhibition • • •

Les Lesson son 3 – Covalent Bonding, Val Valence ence Bond and Molecular O Orbital rbital Theories Valence Bond Theory = Describes how half-filled atomic orbitals on 2 atoms overlap to create a bond containing paired electrons Molecular Orbital Theory • Molecular Orbital (MO) = The region in space where an electron is likely to be found • Bonding Molecular Orbital = s orbitals combine constructivel constructivelyy and causes electron density between atoms to increase (high probability of electrons being found between atoms => MO contributes to bonding) • Antibonding Molecular Orbital = s orbitals combine destructively and causes electron density between atoms to decrease (0 probability of electrons being found between atoms => MO negates from bonding) • When 2 atomic orbitals combine, the bonding MO is lower in energy than the atomic o orbitals rbitals and the antibonding MO is higher in energy • Total Number of Molecular Orbitals = Number of Contributing Atomic Orbitals • Only orbitals of the same symmetry will combine •

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How are MOs fformed ormed from p orbitals? o If pz orbitals overlap head on, the px and py orbitals can only overlap side on and hence they o

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MOs for First Row D Diatomic iatomic Molecu Molecules les o e.g. F2 (14 valence electrons)

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form  MOs When forming MOs, the pz by convention undergoes  bonding and the px and py undergo  bonding

2pz orbitals overlap to give  and  MOs 2px and 2py orbitals overlap to give  and  MOs B2 to N2 have a different diagram The changing order of the MOs can be traced to the increasing separation of the 2s and 2p orbitals as we go across the period Orbitals can only overlap if they are close in energy – there is a big separation between the s orbitals and the p orbitals in O2 and therefore cannot overlap However, in B2 to N2, there is some contribution from the s and p z orbitals to all the  orbitals, whereas from O onwards the are more purely p-like or s-like – this is why there is a reverse in the position of the  MOs and the  MOs

Disadvantage of the MO Theory = Difficult to predict the order in energy of orbitals except for the simplest molecules Higher level computation and comparison to experimental data is required

• Liquid oxygen is found to be paramagnetic – it must possess unpaired electrons Bond Order = [(No. of Bonding Electrons) – (No. of Antibonding Electrons)] Electrons)]/2 /2

Section B – Spectroscopy Les Lesson son 1 – Light

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Electromagnetic Theory of Light = Electrical fluctuations described by the electric field strength E, and the magnetic fluctuations described by the magnetic field intensity (B or H), travelled through space in unison at a constant speed c = 3 x 108 ms-1

 = Wavelength (m) v or f = Frequency (s-1 or Hz) Infrared Spectroscopy = Detects vibrations between atoms in molecules – used to identify molecular structure and functional groups • Nuclear Magnetic Resonance (NMR) = Uses light in the radiowave region to study spin flips of atomic nuclei in magnetic fields and gain detailed information about local structural environment • UV UV-Visibl -Visibl -Visible e = Electron transitions between quantized energy levels in atoms and molecules – used to study bonding • Light is both a particle and a wave o Newton thought of light as a particle (photon) with properties of mass (m) and momentum (mv) o v = f. f.  where v = speed and f = frequency o WaveWave-particle particle Duality Equation -  = h/mv where h, Planck constant = 6.626 x 10-34Js-1, m = mass and v = frequency o E = hv = hc/ where E = energy of a photon o Higher energy radiation is light with high frequency and short wavelength Bonding, Electronic Energy Levels and Transitions Between Them • • •

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The electron energy is related to the radius of the orbit This is determined by the principal quantum number Electrons can jump to a higher energy orbit by absorbing light in the UV to visible range

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The energy jump (E) is related to the frequency or wavelength of the light absorbed by Planck’s equation

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The higher the frequency of light absorbed, the larger the transition Molecular Orbitals of C=C Bonds o Consider the  −  transition, between the HOMO and LUMO energy levels o

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The easiest transition of electrons is from the  −  energy level as it takes the least amount of energy to jump

o The absorbance shown on the graph at 178nm ( max) corresponds to the  −  transition Wavenumber Unit = 1/ in cm-1

Les Lesson son 2 – UV and Visible Spectro Spectroscopy scopy •

Relating Visible Absorption Spectra to Colour = The visible colour is opposite the absorbed colour

 - * Transitions in Polyunsaturated Hydrocarbons with Conjugated -C-C -C-C=C=C- Bonds • The  - * orbitals interact with each other and mix to create a new pair of 1/2 and *1/*2 MOs and energy levels, that are both equally spaced around the positions of the constituent orbitals • This means that the spacing between the highest occupied energy level and the lowest unoccupied level becomes smaller • The electrons derived from the participating ethylene units fill into the 1 and 2 levels, so that the lowest energy transition occurs from 2 into *1 Longer Chains • As we add more double bonds to the conjugated polyunsaturated system, the process repeats, so that more  and * orbitals mix together and the spacing betwee between n the HOMO and LUMO becomes smaller as the chain gets longer • This means that the - HOMO-LUMO transition shifts to longer wavelength as the chain length increases • On average, for longer chain hydrocarbons, the shift is ~30nm for each C=C bond unit we have added to the initial ethene unit • To estimate the effect on the  - * absorption wavelength – 30n + 180 – where n = number of C=C bonds added to the ethene  - * Transitions in Aromatic Ring Systems – Resonance Hybrid of C=C and C C-C -C Bonding (Benzene) • Overlap of pz orbitals to form a , * system Each sp2 hybridised C atom has 4 valence electrons 1 electron points out of the ring to form a  bond with H 2 electrons form  bonds with C atoms on each side The remaining p electrons form a  system with 3 (occupied) and 3* (empty) orbitals above and below the ring Polyc Polycycli ycli yclicc Aromatic Ring Systems -  - * Transitions • • • •

As for the conjugated hydrocarbons, max moves to longer wavelength as the number of aromatic rings involved increases Chlorop Chlorophyll hyll a • The 2 main absorption peaks are 430nm and 660nm • This means that light in the blue-violet and red-orange regions is absorbed, so that compounds containing chlorophyll appear green-yellow when illuminated by white light • This explains the colour of leaves and certain plant stems Haemoglobin in RBCs • The red colour of the haem group at the centre of the protein complex arises from the aromatic  - * transitions of the surrounding porphyrin ring •

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The colour is modified by changes in the electronic state of Fe 2+ as it changes between low-spin and high-spin states as the haem group absorbs O2 in the lungs and delivers it to muscle tissue and other organs throughout the body

• Results in the difference in colour between arterial and venous blood Beer-L Beer-Lambert ambert Law •

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The Beer-Lambert relation derives from the fact that if you send a beam of light through a system of absorbing molecules or centers, then its intensity will decrease exponentially as it travels through the sample, according to the wavelengths of absorption by the sample Absorbance =  c.d c.d = log10(Io/It) = –log10T o  = molar absorption or extinction coefficient, with units Lmol-1cm-1 o c = concentration in mol/L o d = path length in cm

Les Lesson son 3 – Infrared Spectro Spectroscopy scopy

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Infrared spectroscopy measure the vibrational frequencies of molecules For non-diatomic molecules, the different atomic vibrations are organised into patterns, called vibrat vibrational ional modes The frequency at which the vibrations occur is determined by: o Masses of the atoms involved o Strength of the bond For a molecule with n atoms, there are 3n – 6 vibrational modes

• For linear molecules , the number of vibrations is 3n – 5 Quantify Quantifying ing Vibrations • To determine the vibrational frequenc frequencyy (v0, measured in s-1), we use this equation

k = force constant – reflects the resistance of the bond to the extension or compression and is related to bond str strength ength  = reduce reduced d mass – introduced to remove molecular translations and rotations from the equation so that vibrational motions are considered only

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The force constant of a bond indicates its stiff stiffness ness The greater the force constant, the greater the stiffness of tthe he bond A higher vibrational frequency means we are dealing with a stiffer bond and usually a higher bond strength

Calculation Inst Instructions ructions • Convert  into SI units of kg – multiply by the rest mass of a photon (1.67 x 10-27kg) • •

To convert v0 into wavenumber units, divide by the speed of light (c = 3 x 1010cm/s) To obtain v0 directly in wavenumber units, we need to use the reduced mass in atomic mass units and use the equation

Isotope Effe Effect ct – Reduced Mass and Vibrational Frequencies

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Different isotopes of the same element have the same number of protons and electrons but different number of neutrons in the nucleus

We can use isotopic substitution to identify which vibrations involve a particular atom or group of atoms in the IR spectrum, based on whether or not we observe an isotopic frequency shift IR Spe Spectroscopy ctroscopy • IR absorption occurs when there is a change in molecular dipole moment during vibration •

• Vibrational oscillations occur with frequencies v = 1012 – 1014s-1 IR A Activity ctivity of Dif Differ fer ferent ent Molecula Molecularr Vibrat Vibrations ions • Homonuclear Diatomic Molecules o No dipole moment o No change in reduced mass during the vibrations o Therefore, does not show IR absorption • Heteronuclear Diatomic Molecules o Have dipole moment that changes during vibration o Shows IR absorption • Polyatomic Molecules o Sometimes some vibrations do not show IR absorption because of symmetry

Vibra Vibrational tional Modes • • • •

Descriptive labels – stretching, bending, wagging, torsional or deformation Designated as symmetric or asymmetric Asymmetric vibrations occur at higher frequency than symmetric vibrations Bending vibrations usually occur at about half the frequency of stretching modes

Les Lesson son 4 – NMR Spectroscopy Atomic nuclei are charged particles spinning around the nuclear axis which causes them to give rise to a local magnetic field • In an external magnetic field (BO), the axis tends to become aligned with or against the external fiel d direction • Most nuclear spins tend to align with the field because it is more energetically favourable The Larmor Frequency (vL) • The nuclear spin axis is tilted with respect to the external field and precesses around it , like a spinning top • The frequency of precession is known as the Larmor frequency •

• Where  = gyromagnetic ratio and BO = strength of the external applied magnetic field • The Larmor frequency occurs in the radiofrequency range NMR Spec Spectroscopy troscopy Mechanics • NMR spectroscopy occurs by absorption of radiofrequency (RF) radiation by causing the spins to reverse their up/down direction relative to the external field axis • We place a molecule containing atoms with spin-active nuclei in an external magnetic field of known BO, then send in electromagnetic radiation with frequencies tuned within the radiofrequency range • We then record the signal by a radio receiver and look for frequencies where absorption occurs Quantum Mechani Mechanical cal Model of NMR • Both protons and neutrons contribute to determining the spin quantum number (I) • The number of quantised spin states is 2I + 1 – so a ½ spin nucleus has 2 possible spin states (up or down) • If the number of neutrons + protons is odd, then the nucleus has a half integer spin • If the number of neutrons and protons are both odd, then the nucleus has an integer spin • When there is no field present, both +1/2 and -1/2 states have the same energy • But when an external magnetic field is applied, the different spin states have different energies, causing a separation between the 2 energy levels • The separation gets larger as BO is increased • E = h.vL NMR Experime Experiment nt • A high field magnet is used to provide the external field – a flask around the magnet is filled with liquid nitrogen and liquid helium to cool down the superconducting magnet • A radio transmitter sends the electromagnetic signal to the sample to cause the spins to flip • A radio receiver records the signal • We tune the coil and receiver to the characteristic Larmor frequency range of the nucleus we want to observe • The sample is placed as a liquid or dissolved in a solvent in an NMR tube • We also add a standard (TMS) to calibrate the NMR signal • We spin the sample gently to make sure the chemical environments are averaged in the field

Shieldi Shielding ng and Chemical Shift • Nuclei in molecules are surrounded by electron clouds • The circulating electrons cause a small magnetic field that opposes the external applied magnetic field, BO • Shielding = The actual magnetic field felt by the nucleus is slightly lower than expected • The effective field felt by the nucleus is – Beff = BO(1 - ) where  is the shielding constant • The result is a chemical shift in the NMR frequency • More shielding leads to a lower field while less shielding leads to higher field at the nucleus • If the nucleus is more de-shielded, it experiences a higher local field which means that the difference in energy between the spin up and spin down states is larger – so frequency increases Shieldi Shielding, ng, Electronegativity and Chemical Shift • If the adjacent atom is electron withdrawing, it decreases the electron density around the proton under investigation • The amount of shielding then decreases and its NMR frequency increases NMR Standards and Chemical Shift Scale ( values) • To compare the chemical shift values obtained in different labs or on different instruments reliably, a reference is used

The usual standard for NMR is TMS (tetramethylsilane) because o TMS is unreactive o Liquid and mixes with common solvents o Has a single NMR peak that is easily recognisable and occurs at a different position to protons in most organic compounds o Most  values are +ve beca...