2 Rutherford’s Gold Foil Expt. ObservationInferenceAtom is mostly empty space.Positively charged part of atom must be small.Core of atom is very small, dense and positively charged.Most α particles pass through undeflected.Some α particles are slightly deflected.Few particles are reflected back (i.e. do not pass through).
3 Bohr Model and Emission Line Spectra of Hydrogen Observation4 distinct bands of colourInferencee- can only posses certain energies (quantized)Certain wavelengths of light are absorbed by atom as e- transitions from lower to higher energy level (orbit).As e- returns to lower energy level, photons of light are released.If energy released corresponds to a wavelength of visible light, a specific band of colour will be observed.
4 Quantum Mechanical Model Niels Bohr &Albert EinsteinModern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).
5 Electrons as WavesEnergy is quantized. photons –> particles of light(E=hν Planck)Matter has wavelike properties. (deBroglie)Louis de Broglie~1924Standing wavesn = 4Standing Wave20015010050- 50-100-150-20020015010050- 50-100-150-200Forbiddenn = 3.3Adapted from work by Christy Johannesson
6 Schrodinger’s Wave Equation Used to determine the probability of finding the electron at any given distance from the nucleusOrbital “electron cloud”-region in space where the electron is likely to be found90% probability offinding the electronOrbital
7 Heisenberg Uncertainty Principle p. 200 Crazy Enough?Heisenberg Uncertainty PrincipleImpossible to know both the velocity and position of an electron at the same timeWerner Heisenberg~1926gMicroscopeIn order to observe an electron, one would need to hit it with photons having a very short wavelength.Short wavelength photons would have a high frequency and a great deal of energy.If one were to hit an electron, it would cause the motion and the speed of the electron to change.Lower energy photons would have a smaller effect but would not give precise information.Electron
8 Modern View The atom is mostly empty space Two regions Nucleus protons and neutronsElectron cloudregion where you might find an electron
9 Models of the Atome++-Dalton’s model(1803)Greek model(400 B.C.)Thomson’s plum-puddingmodel (1897)Rutherford’s model(1909)Bohr’s model(1913)Charge-cloud model(present)1803 John Daltonpictures atoms astiny, indestructibleparticles, with nointernal structure.1897 J.J. Thomson, a Britishscientist, discovers the electron,leading to his "plum-pudding"model. He pictures electronsembedded in a sphere ofpositive electric charge.1911 New ZealanderErnest Rutherford statesthat an atom has a dense,positively charged nucleus.Electrons move randomly inthe space around the nucleus.1926 Erwin Schrödingerdevelops mathematicalequations to describe themotion of electrons inatoms. His work leads tothe electron cloud model.1913 In Niels Bohr'smodel, the electrons movein spherical orbits at fixeddistances from the nucleus.1924 Frenchman Louisde Broglie proposes thatmoving particles like electronshave some properties of waves.Within a few years evidence iscollected to support his idea.1932 JamesChadwick, a Britishphysicist, confirms theexistence of neutrons,which have no charge.Atomic nuclei containneutrons and positivelycharged protons.1904 Hantaro Nagaoka, aJapanese physicist, suggeststhat an atom has a centralnucleus. Electrons move inorbits like the rings around Saturn.Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
10 Homework Describe the current model of the atom. Describe the contributions of deBroglie, Schrodinger and Heisenberg to the development of the modern view of the atom.
11 ΔE is released as a photon of light. Explaining emission spectrum:ΔE is released as a photon of light.
12 Sommerfeld (1915) -main lines of H spectrum are actually composed of more lines -introduced the idea of sublevels which is linked to the orbital shape
14 Zeeman effect - Single lines split into multiple lines in the presence of a magnetic field. Orbits can exist at various angles and their energies are different when the atom is near a strong magnet.
15 Stern-Gerlach experiment – a beam of neutral atoms is separated into two groups by passing them through a magnetic field
16 Orbitals s px pz py carbon y z x 2s 2p (x, y, z) Using Computational Chemistry to Explore Concepts in General ChemistryMark Wirtz, Edward Ehrat, David L. Cedeno*Department of Chemistry, Illinois State University, Box 4160, Normal, ILMark Wirtz, Edward Ehrat, David L. Cedeno*
17 Y21sY22sY23srrrrrrDistance from nucleus(a) 1s (b) 2s (c) 3s
18 p-OrbitalspxpzpyZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 335
25 Principal Energy Levels 1 and 2 Note that p-orbital(s) have more energy than an s-orbital.
26 Homework Describe the shape of s and p orbitals. Define orbital. How does the size of the orbital change as n increases?
27 Quantum Numbers (used to represent different energy states) Pauli Exclusion PrincipleNo two electrons in an atom can have the same 4 quantum numbers.Wolfgang Pauli1. Principal # (n) 2. Secondary # (l) 3. Magnetic # (ml) 4. Spin # (ms) energy levelsublevel (s,p,d,f), shapetilt or orientation of orbitalelectron spinKey expt workl hydrogen spectral lines made up of closely spaced linesml splitting of spectral lines caused by a magnetic fieldms magnetismCourtesy Christy Johannesson
28 Quantum Numbers 2s 2px 2py 2pz l = 0 (s) l = 1 (p) l = 2 (d) l = 3 (f) Courtesy Christy Johannesson
29 Quantum Numbers n (principal quantum #) n = 1,2,3… l (secondary quantum #)l = 0,1…n-1ml (magnetic quantum #)ml = -l to +lms (spin quantum #)ms = ½, -½
30 Quantum Numbers n = # of sublevels per level Principalleveln = 1n = 2n = 3SublevelsspspdOrbitalpxpypzpxpypzdxydxzdyzdz2dx2- y2An abbreviated system with lowercase letters is used to denote the value of l for a particular subshell or orbital:l =Designation s p d f• The principal quantum number is named first, followed by the letter s, p, d, or f.• A 1s orbital has n = 1 and l = 0; a 2p subshell has n = 2 and l = 1(and contains three 2p orbitals, corresponding to ml = –1, 0, and +1); a 3d subshell has n = 3 and l = 2 (and contains five 3d orbitals, corresponding to ml = –2, –1, 0, –1, and +2).n = # of sublevels per leveln2 = # of orbitals per levelSublevel sets: 1 s, 3 p, 5 d, 7 fCourtesy Christy Johannesson
31 Maximum Capacities of Subshells and Principal Shells n nlSubshelldesignation s s p s p d s p d fOrbitals insubshellAn abbreviated system with lowercase letters is used to denote the value of l for a particular subshell or orbital:l =Designation s p d f• The principal quantum number is named first, followed by the letter s, p, d, or f.• A 1s orbital has n = 1 and l = 0; a 2p subshell has n = 2 and l = 1(and contains three 2p orbitals, corresponding to ml = –1, 0, and +1); a 3d subshell has n = 3 and l = 2 (and contains five 3d orbitals, corresponding to ml = –2, –1, 0, –1, and +2).Relationships between the quantum numbers and the number of subshells and orbitals are1. each principal shell contains n subshells;– for n = 1, only a single subshell is possible (1s);for n = 2, there are two subshells (2s and 2p);for n = 3, there are three subshells (3s, 3p, and 3d);2. each subshell contains 2l + 1 orbitals;– this means that all ns subshells contain a single s orbital, all np subshells contain three p orbitals, all nd subshells contain five d orbitals, and all nf subshells contain seven f orbitals.SubshellcapacityPrincipal shellcapacity n2Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320
32 Practice Decide whether each set of quantum #s is allowed. n = 1 l = 0 ml = 0 ms = ½n = 2 l = 1 ml = -1 ms =- ½n = 3 l = 3 ml = 2 ms = ½n = 0 l = 0 ml = 0 ms = ½2. Give the complete set of quantum #s for a lone e- in the third orbital of a 2p sublevel.