Chapter 2 - Historical Development of Atomic Theory Flashcards
What did Dalton discover?
Atoms combine in simple numerical ratios to form compounds
What did Avogadro discover?
Equal volumes of gas at equal temperatures and pressures contain the same number of molecules
What did Thomson discover?
JJ Thomson came up with the plum pudding model (stable thing with “blueberries” in it), which says that an atom has electrons surrounded by a soup of positive charge. He demonstrated that atoms are actually composed of aggregates of charged particles. Prior to his work, it was believed that atoms were the fundamental building blocks of matter. He was able to determine the charge to mass ratio of the electron.
What did Rutherford discover?
Rutherford overturned Thomson’s model in 1911 with his well-known gold foil experiment in which he demonstrated that the atom has a tiny, heavy nucleus. He also came up with the mass of the electron.
The Rutherford model is that the atom is made up of a central charge (this is the modern atomic nucleus, though Rutherford did not use the term “nucleus” in his paper) surrounded by a cloud of (presumably) orbiting electrons.
What did Canizzaro discover?
In 1860, he came up with a consistent set of atomic weights. Because each molecule is made up of atoms, the sum of the atomic weights can then give us the molecular weight.
What did Mendeleev + Meyer discover?
They came up with the periodic table. They organized the periodic table by similar properties.
Period
rows in periodic table
Group
columns, also known as “families” in the periodic table
IUPAC vs US groups?
IUPAC: groups numbered 1–18
US: groups IA–VIII main group, “B” group are transition metals
Balmer
Balmer showed that energies of light emitted by the hydrogen atom are given by the equation:
E = R ( 1/2^2 - 1/n^2) = R ( 1/nl^2 - 1/nh^2)
where R = Rydberg constant = 1.097E7 m-1 = 2.179E-18 J
Energy related to wavelength, frequency, wavenumber
E = hv = hc/λ
de Broglie
λ = h/mv
Heisenberg Uncertainty principle
there is a relationship between the inherent uncertainties in the location and momentum of an electron.
∆x∆p ≥ h/4π
Therefore, since we cannot know the position exactly, we cannot use ORBITS but must use ORBITALS, regions to descibre the probable location of electrons.
Electron density
the probability of finding an electron at a particular point in space
What are some problems with the Bohr model of the atom?
- works well only when one electron is involved
- spherical –> elliptical orbitals introduced to fit Bohr’s data
- did not take into the wave-like nature of the electron (de Broglie)
- Heisenberg uncertainty principle means that we cannot know the position of the electron exactly. Therefore, we cannot use orbits, but must use orbitals instead to describe the probable location of electrons.
Schrodinger Equation
- tells us the wave properties of an electron, its mass, position, total energy, and potential energy
- based on the wavefunction ψ, which describes the wave properties of a given electron in a particular orbital
Hψ = Eψ
What is the relationship between atomic orbitals and ψ?
Atomic orbitals are described by a unique ψ, so there is no limit to the number of solutions for the Schrödinger equation.
Bohr’s Quantum Theory of the Atom
Negatively charged electrons in atoms moves in stable circular orbits around the positively charge nucleus with no absorption or emission of energy.
The energy emitted/absorbed by electron can be found by:
E = R ( 1/nl^2 - 1/nh^2)
µ in the definition of the R used in the Bohr theory?
Reduced mass of the electron/nucleus combination.
1/µ = 1 / mass of electron + 1/mass of nucleus
L
Angular momentum/shape of the orbital. Describes angular dependence and contributes to energy.L = 0,1,2,…, n-1
M_L
Orientation of the angular momentum vector in a magnetic field. Describes orientation in space (angular momentum in z direction).m_l = -l, …., -2, -1, 0, 1, 2, …, +l
M_s
Spin. Orientation of the electron’s magnetic moment (spin) in a magnetic field, either in direction of the field (+1/2) or opposed to it (-1/2).m_s = -1/2, +1/2
What l value describes:1. S2. P3. D4. F5. G
- 4
N
Principal quantum number (n) is primarily responsible for determining the overall energy of an atomic orbitaln = 1,2,3,…
Radial functions
the radial part of the wavefunction: electron density at different distances (r) from the nucleus
- determined by n and l
- radial prodbability function = 3πr^2 R^2
Radial nodes
radial nodes for an orbital = n - l - 1
when R, the radial function, = 0
(counts a conical node surface such as for a d orbital as two nodes)
Angular functions
the angular part, Y(ϴ,ϕ). describes the shape of orbitals and their orientation in space
- determined by l and ml
Angular nodes
when Y=0, are planar/conical
Nodal surface
A surface with zero electron density
Node
A surface where the wavefunction is zero as it changes sign
ψ=0 so either R=0 or Y=0
Total number of nodes in any orbital?
n-1
Total number of radial nodes for any orbital?
n-l-1
Total number of angular nodes for any orbital?
l
Aufbau principle
Limitations on the values of quantum numbers –> Aufbau (building up) principle, where the buildup of electrons in atoms results from continuously increasing quantum numbers
- Electrons placed in give lowest total energy
- Pauli exclusion principle (each electron has unique set of quantum numbers)
- Hund’s rule of maximum multiplicity (electrons placed to have maximum total spin)
Pauli Exclusion principle
Each electron has a unique set of quantum numbers
Hund’s rule
Electrons placed to have maximum total spin.
- comes from the coulombic energy of repulsion (πc) vs the exchange energy (πe)
Coulombic energy of repulsion
destabilizing effect caused from electron repulsion between electrons in the same orbital
Exchange energy
stabilizing effect of increased number of possible exchanges, a consequence of parallel electron spins in separate orbitals.
Klechkowsky’s rule
The order of filling of the orbitals proceeds from the lowest available value for the sum n+l
Shielding
each electron can act as a “shield” for electrons farther from the nucleus, reducing the attraction between the nucleus and the more distant electrons
Z* = Z - S
where Z* = effective nuclear charge
S = sheilding constant, determined using Slater’s rules
Slater’s rules
rules for determining the shielding constant for a specific electron
For transition metal atoms, which electrons will be lost more readily?
ns electrons are lost more readily than (n-1)d electrons. This is because the effective nuclear charge for the ns electrons are smaller than the values for the (n-1)d electrons, so they are held less tightly.
For Cu and Cr, what is the exception for shielding?
Half and fully filled d and f shells are preferred, so Cr has a configuration of [Ar]4s13d5 instead of [Ar]4s2d4.
Ionization energy
Energy required to remove an electron from a a gaseous atom/ion
- generally increases across a period
- break at Boron due to shielding
Electron affinity
Energy required to remove an electron from a negative ion (zeroth ionization energy)
What has the lowest electron affinities?
Noble gases (an extra electron beyond the noble gas configuration is easily lost)
Anions are ____ than cations with similar numbers of electrons.
Larger (e.g. electron repulsion)
Covalent radii/atomic size changes?
Increasing nuclear charge + # electrons –> decrease in atomic size across period
