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ELECTRONSTopic 1B
THE ELECTRON J.J. Thomson, English (1898-1903)—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-
COULOMBS LAW 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 108 C/g e is charge on electron in Coulombs, (C) and m is its
mass. Or F=q1q2/r2
If q1 and q2 have the same sign the force is repulsive, when opposite the forces is attractive
BOHR MODEL OF THE ATOM
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 & THE SPECTROSCOPE 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)
EMISSION SPECTRA 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??
LYMAN, BALMER & PASCHEN
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
EXAMPLES OF EMISSION (LINE) SPECTRA
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!!
IONIZATION ENERGY
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.
TYPES OF SPECTRUM
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
ABSORPTION OF LIGHT
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
ELECTRONIC CONFIGURATION
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.
FROM PLANETARY TO QUANTUM MECHANICAL MODEL
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
ARRANGEMENT OF ELECTRONS IN ATOMS
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
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, ms.
ARRANGEMENT OF ELECTRONS IN ATOMSARRANGEMENT OF ELECTRONS IN ATOMS
ELECTRON SPIN QUANTUM NUMBER, MS
Can be proven experimentally that electronhas a spin. Two spin directions are given byms where ms = +1/2 and -1/2.
Can be proven experimentally that electronhas a spin. Two spin directions are given byms where ms = +1/2 and -1/2.
ELECTRON SPIN QUANTUM NUMBER
Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.
Diamagnetic: NOT attracted to a magnetic fieldParamagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.
n ---> shell 1, 2, 3, 4, ...
l ---> subshell 0, 1, 2, ... n - 1
ml ---> orbital -l ... 0 ... +l
ms ---> electron spin+1/2 and -1/2
n ---> shell 1, 2, 3, 4, ...
l ---> subshell 0, 1, 2, ... n - 1
ml ---> orbital -l ... 0 ... +l
ms ---> 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.
QUANTUM NUMBERS & ATOMIC ORBITALS 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.
ANGULAR MOMENTUM 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’
MAGNETIC QUANTUM NUMBER 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 CONFIGURATIONSWRITING ATOMIC ELECTRON CONFIGURATIONS
11 s
value of nvalue of l
no. ofelectrons
spdf notationfor H, atomic number = 1
Two ways of writing configs. One is called the spdf notation.
Two ways of writing configs. One is called the spdf notation.
WRITING ATOMIC ELECTRON CONFIGURATIONSWRITING ATOMIC ELECTRON CONFIGURATIONS
Two ways of writing configs. Other is called the orbital box notation.
Two ways of writing configs. Other is called the orbital box notation.
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
One electron has n = 1, l = 0, ml = 0, ms = + 1/2Other electron has n = 1, l = 0, ml = 0, ms = - 1/2