TOPIC 2 Flashcards
Atomic Theories
Dalton
1803
Thompson
1904
Rutherford
1911
Bohr
1913
Schrödinger
1926
Rutherfords gold foil experiment
Rutherford recalled that “it was as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you”
He shot a beam of alpha particles coming from a radioactive source at a sheet of gold foil, and a with a detector identified that some of the particles were deflected by small angles, indicating a small concentrated positive charge; he concluded that a small, dense nucleus (positively charged) was causing the deflections.
Mass number and Atomic number
Atoms of and element are characterized by 2 numbers
A = Mass Number = p+ + n• = number of nucleones (particles in the nucleus of the atom)
Mass number A can only be given for an atom of an element (not an element as a whole) and must always be a whole number.
Z = Atomic Number = p+ = position of an element in the P.T. = nuclear charge
Atoms are…
Neutral.
This means that they have the same amount of p+ and e-
Relative mass and charge of e-, p+ and n•
p+ n• e-
Relative mass 1 1 0
Charge +1 0 -1
No unit, its mass in comparison with themselves.
The e- has mass, because its matter, but in comparison to the p+ and n• its nearly neutral.
Ions
Ions are formed when the number of p+ in an atom is no longer balanced by the number of e- in the atom.
They are chemically charged elements
When an atom loses e-, a + ion is formed
- cation
When an atom gains e-, a - ion is formed
- anion
Isotopes
Atoms of the same element with different mass number
They have the same atomic number, and the same chemical properties, but different number of n• and different physical properties.
Mass number (Ar)
Mass number is called the relative atomic mass of an element (Ar).
The Ar of an element takes into account the relative abundances of its isotopes.
“The Ar of an element is the average of its isotopes, considering their abundances.”
Ar = ?
Ar = A x %abundance / 100
A = mass number
Schrödinger’s model of the atom
He introduces atomic orbitals: an atomic orbital is a region around the nucleus in which we are most likely to find and e- (95% - 99%)
ORBIT AND ORBITAL ARE DIFFERENT
(Orbit is a fixed path [main energy level’s - Bohr’s model of the atom])
Schrödinger stated that e- do not move in set paths around the nucleus, but in waves. Its impossible to know the exact location of an e-; instead, we have “clouds of probability ” calles orbitals, in which we are more likely to find e-
Atomic orbitals characteristics
Any orbital can hold a maximum of two e- of opposite spin.
Atomic orbitals are found within a specific energy sub-level, within a main energy level.
Main and Sub energy levels
Types of sub-levels
S: 1 orbital: 2e- max
P: 3 orbitals: 6e- max
D: 5 orbitals: 10e- max
F: 7 orbitals: 14e- max
Main energy levels
1) 2e- max
s: 2e-
2)
8e- max
s: 2e-
p: 6e-
3) 18e- max
s: 2e-
p: 6e-
d: 10e-
4) 32e- max
s: 2e-
p: 6e-
d: 10e-
f: 14e-
5)
s
p
d
f
n)
…
For the same main energy level ‘n’, the energy of the sub-levels increases as s<p<d<f
“Rule of rain”
Oder in which we fill the orbitals and sub-levels
(Tabla)
Num: main energy level (Bohr)
Letter: energy sub-levels (Schrödinger)
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
Increasing energy—>
As the e- jumps away from the nucleus, the electromagnetic force of attraction towards it decreases. As a result, it become less stable, therefore accounting for a higher energy.
e- configuration
Full ground state (normal e- configuration- no excited e-)
Condensed (use the noble gas previous to the element to reduce the configuration, then continue until reaching the element on the P.T.)
e- configuration of ions
When + ions are formed, we first remove e- with the highest energy (highest main energy level)
The e- configuration of - ions is determined by simply adding the e- into the next available electron orbital.
Orbital diagrams
Atomic orbitals are represented as “boxes”, each holding a maximum of 2e- of opposite spin.
The spin of the e- is indicated via an arrow. Each half-arrow represents an e-
s: 1 box
p: 3 boxes
d: 5 boxes
f: 7 boxes
Rules for filling in atomic orbitals
1) Aufball principle: e- are placed into orbitals of lowest energy level first
2) Pauli exclusion principle: two e- cannot occupy the same orbitals unless they have opposite spins.
3) Hund’s rule: when filing orbitals of equal energy (same main and sub energy level), the e- fill the orbitals singly first to avoid repulsion.
e- configuration exceptions
Exceptions:
Cu: [Ar] 4s1 3d10
Cr: [Ar] 4s1 3d5
These exceptions are a consequence of the seem of stability from Cr and Cu.
Cr: the configuration 4s1 3d5 is better (more stable) than 4s2 3d4, as it foes not have e- — e- repulsion from 4s.
Cu: the configuration 4s1 3d10 is better (more stable) than 4s2 3d9 as the sub-level 3d is more attracted to the nucleus (4s is farther from the nucleus than 3d).
Degenerate orbitals
Orbitals of equal energy (same main and sub energy levels)
2px, 2py, 2pz
Shapes of orbitals
s
O
p
(Infinito)
px (infinito horizontal)
py (infinito vertical)
pz (sale de dimensión/ para ‘adelante’)
Spectra
Electromagnetic (EM) radiation comes in different forms of differing energy. Together they make up the EM spectrum.
R.w M.w IR [visible] UV X-rays Gama-rays
————————————————> E, f
Wavelength <————————————
Energy is directly proportional to frecuency
Energy is inversely proportional to wavelength.
Continuous and discrete spectrum
The continuous spectrum is the visible region of the EM spectrum. It contains all the colors which together make up white light (w.l. is a mixture of light waves of differing wavelengths or colors)
The discrete/line spectrum is characteristic to each element and only has certain light waves present (only certain colors will appear when the photon emitted by the e- of that element falls into the visible region of the EM spectrum)
e- are found in discrete (particular/specific energy) energy levels outside the atoms nucleus
In the context of the atomic spectra we will be always working with Bohr’s model of the atom.
e- changing energy levels
e- can move to a higher energy level (excited state) by absorption of a photon. Similarily, e- can move from an excited state to a lower energy level by emitting a photon.
If the photon emitted by an e- as it falls to a lower energy level has an energy wave within the visible region of the EM spectrum, we will interpret it as a specific color.
Photon: a particle of light (dual particle-wave behavior of light)
Bohr model of the atom and why we use it in atomic spectra
In Bohr’s model of the atom, energy levels are discrete and converge towards increasing energies.
(Drawing)
In atomic spectra we work with Bohr’s model of the atom because e- transitions between energy sub-levels (Schrödinger) can not be distinguished with precision, at least for now.
Line spectra are evidence of e- being found on discrete energy levels. If this were not true, the absorption spectrum would consist in all black, and the emission spectrum would be the same as the continuous spectrum.
e- transitions in an H atom
Emission Absorption
nf radiation ni radiation
1 UV 1 UV
2 vis 2 vis
3 IR 3 IR
Valence e-
The total number of e- in the outer main energy level (we must add up all the e- located in the outer main energy level, independently from their sub-level)
The group number of an element indicates its valence e-
Example: Be belongs to group 17 —> 7 valence e-
Emission and Absorption spectrums
Emission spectrum:
When an e- moved from higher to lower energy levels.
Absorption spectrum:
When and e- moves from lower to higher energy levels.
Ground state
Energy level an e- normally occupies (not excited)
Wavelengths
“Longer wavelength” = less energy
“Shorter wavelength” = more energy
