Topic 1B

 J.J. Thomson, English ( )—found that when high voltage was applied to an evacuated tube, a “ray” he called a cathode ray [since it emanated from the (-) electrode or cathode when YOU apply a voltage across it] was produced.  The ray was produced at the (-) electrode  Repelled by the (-) pole of an applied electric field, E  He postulated the ray was a stream of NEGATIVE particles now called electrons, e-

 He then measured the deflection of beams of e- to determine the charge-to-mass ratio  e is charge on electron in Coulombs, (C) and m is its mass.  The interaction between charged objects is a non- contact force that acts over some distance of separation. Charge, charge and distance  e/m =-1.76 x 10 8 C/g  e is charge on electron in Coulombs, (C) and m is its mass.  Or F=q1q2/r 2  If q1 and q2 have the same sign the force is repulsive, when opposite the forces is attractive

 Neils Bohr originally came up with the idea of the atom containing an dense nucleus being orbited by negatively charged electrons.  Bohr’s idea (the planetary model) suggested that electrons could only travel in fixed orbits or shells around the nucleus.

 Bohr used a spectroscope and found that gaseous elements emitted electromagnetic radiation when heated  The light that was emitted in quanta (energy packets) gave a unique pattern of radiation.  It was discovered that the release of radiation was caused by electrons in the atom absorbing energy and being promoted to a shell further way from the nucleus (ie orbit 1 to orbit 2 )

 When an electron falls back down to the ground state, it releases the energy that it absorbed when it was promoted to the higher energy shell (the exited state).  This release of energy creates a line spectrum.  In other words, the energy emitted is observed as light - this is the source of the lines we see in the emission spectra Remember fireworks??

 The Lyman series (red lines) is not visible, in the UV range (n>1 to n=1)  Balmer series is in the visible range (green lines) n>2 to n=2  Others series like Paschen, Brackett and Pfund) lie in the infrared region

Hydrogen: Helium: Carbon: The wavelengths of light that are emitted are characteristic for the particular elements that are present in the gas. Thus the emission (bright line) spectrum is a fingerprint of the gas (or gases) that is (are) emitting the light. This is because the shells are in fixed positions, the difference in energies between them (and the wavelength of the line in its spectra) is also fixed. Video!!

 Tie this to ionization energy – sometimes and electron may gain enough energy to completely over the attraction of the nucleus and can be ejected by the atom.  The energy required to do this is different for each electron  all electrons are at a different distance from the nucleus (plus shielding)  This is ionization energy!  Bohr’s finding are consistent with the observed pattern of ionization energies.

 emission spectrum—the spectrum of bright lines, bands, or continuous radiation that is provided by a specific emitting substance as it loses energy and returns to its ground state OR the collection of frequencies of light given off by an "excited” electron  absorption spectrum—a graph or display relating how a substance absorbs electromagnetic radiation as a function of wavelength  line spectrum--isolate a thin beam by passing through a slit then a prism or a diffraction grating which sorts into discrete frequencies or lines  Johann Balmer--worked out a mathematical relationship that accounted for the 3 lines of longest wavelength in the visible emission spectrum of H. (red, green and blue lines)  Niels Bohr connected spectra, and the quantum ideas of Einstein and Planck: the single electron of the hydrogen atom could occupy only certain energy states, stationary states

 In photon absorption, an electron goes to a higher energy level.  The photon has to have just the right energy to be absorbed.  Its energy has to be equal to the energy difference between the two levels.  Photons of any other energy are not absorbed.  What goes up must come down. Energy absorbed must eventually be emitted

 Bohr’s planetary model only gave us part of the picture.  The fixed orbits did not explain all of the electrons’ properties.  Enter Debroglie- his work suggested that the electron was not just a particle but also wave like.  Schrodinger further developed the wave idea and was able to make predictions about where an electron might be found in an atom.

 Coupled with the Heisenberg uncertainty principle (remember can’t know the position of the electron and momentum at the same time) Schrodinger’s work led to the quantum mechanical model – the 3d probability of any one electron at any point in time within each of the shells.  These 3D probability maps are known as orbitals.

ARRANGEMENT OF ELECTRONS IN ATOMS Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (m l )

Each orbital can be assigned no more than 2 electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, m s. ARRANGEMENT OF ELECTRONS IN ATOMS

ELECTRON SPIN QUANTUM NUMBER, M S Can be proven experimentally that electron has a spin. Two spin directions are given by m s where m s = +1/2 and -1/2. Can be proven experimentally that electron has a spin. Two spin directions are given by m s where m s = +1/2 and -1/2.

ELECTRON SPIN QUANTUM NUMBER Diamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons.

n ---> shell1, 2, 3, 4,... l ---> subshell0, 1, 2,... n - 1 m l ---> orbital -l l m s ---> electron spin+1/2 and -1/2 n ---> shell1, 2, 3, 4,... l ---> subshell0, 1, 2,... n - 1 m l ---> orbital -l l m s ---> electron spin+1/2 and -1/2 QUANTUMNUMBERSQUANTUMNUMBERS The electrons address!

PAULI EXCLUSION PRINCIPLE No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron in an atom has a unique address of quantum numbers.

 There are 4 quantum numbers that describe the “address” and “spin” of an electron.  The value of n limits the possible values of ℓ, which in turn limit the values of mℓ.  n—principal energy level--1 to infinity. Determines the total energy of the electron.  It indicates the most probable [within 90%] distance of the electron from the nucleus.  It’s a measure of the orbital size or diameter (cloud size). 2n2 electrons may be assigned to a shell. The larger the value of n the larger the cloud  Energy levels closer to the nucleus have lower energy. As n increases, the orbital  becomes larger and the electron spends more time farther from the nucleus. An  increase in n also means that the electron has a higher energy and is therefore less  tightly bound to the nucleus.  It’s simply the Energy level that electron is in.  If it’s a 3s electron, n = 3, if it’s a 4d electron, n = 4, etc.

 ℓ--angular momentum--0,1,2,....(n-1) electrons w/in shell may be grouped into subshells [or  sublevels, same thing!], each characterized by its certain wave shape.  Each ℓ is a different orbital shape or orbital type.  - n limits the values of ℓ to no larger than n-1. Thus, the number of possibilities for ℓ is equal to n.  (English translation: 3 sublevels for 3rd E level, 4 for 4th E level, etc.)  - s,p,d,f sublevels - 0,1,2,3 ℓ-values respectively (So, what do spdf stand for? Sharp, principle, diffuse, fundamental--early days of atomic spectroscopy) You can keep going from ‘g’

 mℓ--magnetic—assign the “blanks” in orbital notation with zero on the middle blank and then -ℓ  that means that the range of orbital's is from +ℓ to - ℓ Designates the orbital (room) where the electron can be found.  Gives the DIRECTION IN SPACE that the orbital takes.  specifies to which orbital within a subshell the electron is assigned. Orbitals in a given subshell differ only in their orientation in space, not in their shape.  Values of ml : from –l, ….0, ….+l.  The middle orbital of a subshell has a value of 0. Orbitals to the left of the middle orbital have negative numbers; to the right, they have + numbers. How to write quantum numbers

WRITING ATOMIC ELECTRON CONFIGURATIONS 1 1 s value of n value of l no. of electrons spdf notation for H, atomic number = 1 Two ways of writing configs. One is called the spdf notation.

WRITING ATOMIC ELECTRON CONFIGURATIONS Two ways of writing configs. Other is called the orbital box notation. One electron has n = 1, l = 0, m l = 0, m s = + 1/2 Other electron has n = 1, l = 0, m l = 0, m s = - 1/2

See “Toolbox” for Electron Configuration tool.