Atomic structure PDF

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Chem 1301A notes ...

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Chapter 1: Atomic structure Wavelength- distance between 2 closest equivalence points on a wave Frequency- number of oscillation cycles at a given point per second Amplitude- the height of the wave Light- stream of photons that behave as particles and as waves – energy can be absorbed or emitted as photons – energy of each photon is determined by its frequency and wavelength and uses plancks constant Light relationships: - higher frequency-higher energy - longer wavelength= lower energy - higher amplitude= higher energy (brighter light) Electromagnetic spectrum: complete range of wavelengths – visible spectrum is btw 400 and 750nm Emission of light - when hydrogen gas atoms in a ground state molecule become excited they dissociate and jump to higher energy levels and as they return back to ground state energy they emit light at certain wavelengths depending on how excited they were - Lymen series (n=1) light emitted by H atoms in UV region - Balmer series (n=2) light emitted in visible spectrum - Paschen series (n=3) light emitted in infrared region Bohr model of hydrogen atom - states e- orbit the nucleus and each orbit quantizes the energy of e- within it - as n increases the energy required to jump higher gets smaller - problems with this model- did not apply to atoms with more than 1 e- (none hydrogen atoms) – did not explain splitting in spectrums – does not explain why e- in n=1 don’t fall onto nucleus – does not explain why only some orbits are allowed so instead we prefer the quantum mechanical model: - treats e- as waves NOT particles - e- do not orbit the nucleus - waves can interact with 1) constructive interference- in phase waves that result in the sum of the 2 waves creating a higher amplitude or 2) destructive- opposite phase waves combine and cancel each other out resulting in no amplitude - the 2 slit experiment proves that light behaves as waves not particles- when shine light through two slits we get multiple light a dark bands from constructive and destructive interference, if operated as particles we would only get two bright bands where the slits were Debroglie proposed light behaves as Both waves and particles - 2 types of waves: 1) travelling- every point on a wave has same amplitude 2) standing- different points on the wave have different amplitudes btw 2 fixed ends - said e- behave as standing waves and can only ascillate at certain frequencies - nodes of standing waves are regions where no e- are found and amplitude is 0 - more nodes increase energy of the wave - # of nodes= n (principle #) -1 - Schrodinger’s equation gives probability of finding an e- 95% of the time in an orbital Quantum numbers - characterizes e- principle number (n)- energy and size of orbital - azithumal (l) – orbital shape assigns s,p,d or f depending on number- l= 0-n-1 – if l=0 (s) l=1(p) l=2(d) l=3(f) - magnetic number (ml) – orientation of the orbital – ranges from –l - +l - spin number (ms)- indicates spin - +1/2 = up spin -1/2= down spin - each orbital holds max 2 e- orbitals with same energy are degenerate  Pauli exclusion principle: no 2 e- in an atom can have same 4 quantum numbers, therefore e- in same orbital will have same n, l, and ml, but must have opposite spins Shapes of atomic orbitals

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orbitals represent areas with high likelihood of finding an eAufbao principle: e- occupy lower energy orbitals first then higher energy orbitals valence e—have highest principle # isoelectronic species – species with same e- configuration 4s e- filled before 3d and also removed before 3d because higher energy –“4s are first in and first out”  Hunds rule: when orbitals are degenerate they must be filled singly first and then paired - exceptions are Cr (Ar 4s1 3d5) and Cu (Ar 4s13d10) which fill their 3d orbitals and have a ½ filled 4s orbital because a fully filled 3d (singly or paired) is more stable than a filled 4s Periodic Trends: ENC - (z) how strong of a pull-on valence e- z= #protons-#valence e- increases up and across pt Atomic radius - size of the atom - decreases up and across - b/c down pt has more valence e- and shells and across enc gets higher so e- are pulled in closer and atom is smaller bond length - multiple bonds (double/triple) are shorter than single - decreases up and across because as atomic radius gets smaller the atoms can pack tighter and bonds are smaller Ionic radius - same as atomic radius as long as atoms have same charge - if atom is a cation, positive charge (less e-) then atom is smaller and if atom is anion negative charge (gained e-) then atom is larger - if all atoms are isoelectronic (same e- configuration) arrange largest as most negative ion and smallest as most positive ion Ionization energy - min energy required to remove an e- from an atoms valence shell - increases up and across because as enc increases pull on e- is stronger and requires more energy and as number of shells decreases don’t have shielding effect so harder to remove e- exceptions: Boron b/c only 1 p orbital e- is easy to remove and Oxygen b/c has only one paired p orbital and fully ½ filled p orbitals are more stable so easy to remove that e- ionization energy required increases successively b/c as atom loses more and more e- hold on remaining ones is stronger E- affinity - change in energy for addition of an e- in gas phase atom - increase up and across - more negative dh for adding e- it is favourable and increases e- affinity and and if dh becomes more positive when adding an e- it is not favoured and lowers e- affinity - column 2 has a lower e- affinity b/c to add another e- requires making a whole new orbital and column 15 because all p orbitals have a single e- in them to add another e- would require pairing and overcoming repulsion forces Electronegativity - ability of atom to hold onto its valence e- increases up and across pt - because as enc increases and atoms gets smaller pull on e- gets stronger

Chapter 2 Ionic and covalent bonding Ionic -

metal and non-metal

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EN difference > 1.9 Complete transfer to atom with higher EN value Anions and cations form ionic bonds b/c – and + charges strongly attract High mp – how high depends of distance between ions and magnitude of the charge- 9smaller atoms pack tighter and therefore increase mp and ions with charges 4+ and 4- would have higher mp than same two ions with charges 2+ and 2-) Covalent bonding - non-metal and non-metal - can be polar or nonpolar - polar: - share e- not equally – EN diff btw 0.5 and 1.9 – atom with lower EN is d+ and atom with higher EN is d – nd e- are pulled to the d- atom - nonpolar – equal sharing of e—EN diff less than 0.5 also includes C-H bonds Lewis structures - show bonds, unpaired e- and charges (d+ or d-) - formal charges (cf) – the diff btw # of e- in lewis model vs # of valence ecalculate cf = (#valence e-) – (# lewis e-) # of lewis is the sum of dashes and dots in structure around an atom gives overall charge on a single atom in a compound octet rule - atoms share valence e- to obtain noble gas config. - Applies to main group elements only - C,N,O,F – 8e- and H-2e- Elements in row 3+ can have more than 8eDrawing Lewis structures 1. count valence e- in compound 2. organize atoms 3. distribute extra e- provide all atoms with full octet and any leftover go on central atom - if not enough e- must make double/triple bonds ( take from atoms with lowest EN value first) - best lewis structures have atoms with least # of cf’s 4. assign cf’s if apply Expanded octets - row 3+ can have more than 8 e- in valence shell - P,S,Cl,As,Se,Br Best Lewis structures - lowest cf’s and consistent with EN values (lone pairs placed of most EN atoms first) - exceptions: e- deficient mols. – compounds with atoms wo full octet – very reactive – ex. Borane BH3 unpaired e- if molecule has odd number of valence e- - AKA free radicals – very reactive –ex. NO and NO2 Resonance strucs - when there are multiple correct structures that differ only by e- position - resonance hybrids represent these mols by delocalizing the cf over 1 atom over all atoms of that element - delocalized charges lower energy of the molecule Bond order - single bonds BO=1 - double BO=2 - triple BO=3 - resonance structures use avg BO to account for bond that can be moved - calc. avg BO of a given atom by # of bonds / # of atoms Line diagrams - solid=covalentbond - no atom=c - H bonded to C not shown

- Non C and H’s attached to non C atoms always shown - Lone pairs always shown VSEPR THEORY - lewis strucs only tell arrangement of e- and atoms - VSEPR predicts 3d struc Pairs of e- occupy valence shells These pairs aka regions of e- density ( they repel eachother to max separation) - lone pairs repel more than bonding - multiple bonds treated as single region AXmEn notation - A= central atom - X= atoms in molecule m = # of atoms (X) bonded to central (A) - E= lone pairs n= # of lone pairs (E) on central atom (A) - m+n= # of regions of e- density - when determining 3d shape lone pairs are not included # of regions e- arrangement 2 linear 180d 3 trig planar 120 4 tetrahedral 109.5 5 trig bypyramidal 90 and 120 6 octahedral 90 - arrangement of e- always the same given # of regions of density but 3d shape can change based on lone pairs and repulsion forces there # eangle hybridization AXE Lone Mol. Shape Picture regions arrangement notation pairs 2 Linear 180 Sp AX2 0 linear 3 Trig planar 120 Sp2 AX3 0 Trig planar bent...