CHEM1611 - Lecture notes ALL PDF

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CHEM1611 Structure of Atoms - History - 1808  J Dalton Atomic Theory o All matter made from atoms (cannot create/destroy) o Atoms of one element can’t change into atoms of another element - 1897  JJ Thompson Cathode Rays o Negatively charged particles o All metals produced by same particles o Electrons - 1909  E Rutherford Nucleus of atom o Mostly empty space with electrons o All positive charge is in the nucleus - 1909  Niel Bohr o electrons in orbit (energy levels) o electrons move orbits  energy is absorbed/emitted  ENERGY CORRESPONDS TO LIGHT OF SPECIFIC ENERGY/FREQUENCY Electromagnetic Radiation - Wavelength o Iambda (λ)  Distance between two adjacent identical points of the wave - Frequency o V o Nu  No. wave crests passing a given point per uni time  Light waves travel at same speed (in vacuum)  Speed of light  C = 2.998x108ms-1  Wavelength and Frequency relation o C = λV o Radiation varies in energy  E.g. Up V = Up oscillation = Up energy  Energy = Planck’s Constant x frequency  E = hv  h = 6.626x10-34J

Atomic Emission Spectra - We see combination of wavelengths o Pass light through prism to see lines Spectrum - Only light of certain energies is emitted - Pattern of lines is unique to elements

o Suggests emission is quantised (discrete amounts) Theory and experiment agree for H - Energy of the hydrogen atom orbits is inversely proportional to the square of the orbit no. - E = -Er(1/n2)Z2 o Er = 2.18x10-18J o Z = atomic no. - As ∆E = Ef - EI o Then: ∆E = -2.18x10-18J (1/n2Final - 1/n2Initial)Z2 •

Learning Outcomes − Recognise the historical context of the Bohr model of the atom. − Be able convert between the wavelength, frequency and energy of light. − Be able to calculate the energy of a hydrogen orbit. − Be able to calculate the atomic emission spectrum of a hydrogen atom.

Lecture 3 Cheeky Equations - De Broglie o Wavelength to momentum  λ=h/mv -

What is an orbital?

o Estimate location of electrons at a given energy level Waves  Mechanical Bound - Boundary conditions o Continuous o Single valued (function) o Multiples of a whole number of half wavelengths Quantum numbers (n must be whole and positive) - How big an orbital is o n = 1,2,3 o As n increases, energy level increases - Shape  Angular Momentum Quantum number o L limits (n-1)  Sub orbitals Quantum no. Orbital shape L=0 L=1 2 3

s  sphere p  dumbbells with space d  x shape f

Magnetic Quantum Numbers (mL) - ml = -L, -L+1… 0… L - e.g.

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These describe the ORIENTATION OF ORBITALS o E.g. xyz axis o P orbitals

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The Spin Quantum Number: (ms)

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Each orbital, described by n, l and ml, may contain max of two electrons, one spin +1/2 and the other -1/2

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Learning Outcomes − Explain the meaning of the orbital quantum numbers, n l ml ms . − Understand the designation of orbitals such as 1s, 3d, 4p, 4f. − Recognise the shapes of s, p and d atomic orbitals. − Determine the number of electrons in an orbital/subshell/shell. Lecture 3  Subshell energy - For hydrogen, subshells have the same energy (degenerate) - All other atoms

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Li Na K

o Subshells have different energy levels Subshell differences in energy levels decrease as n increases o Order: 1s...