Atomic Theory Flashcards
Describe the relationship of wavelength, frequency and energy for electromagnetic radiation (light).
Energy is directly proportional to frequency and inversely proportional to wavelength. These relationships can be described by Plank’s equation: E = hv = hc/lambda
In what region of the electromagnetic spectrum (in terms of wavelength) do we find the infrared, visible, and ultraviolet regions?
Infrared is considered to be above 700 nm, visible between 700-400 nm, and finally ultraviolet below 400 nm.
Describe the Rutherford model of the hydrogen atom circa 1910.
Based on the experimental evidence of the time (including Rutherford’s gold foil experiments), the hydrogen atom was understood to consist of a positively charged particle (designated as the nucleus) surrounded by an electron in motion at a relatively large distance from the nucleus. Evidence suggested that most of the mass (99.9%) was associated with the nucleus. From this model, it was inferred that strangely, the atom and therefore matter consisted of mainly empty space.
According to classical physics, what is the major flaw of the Rutherford model, which renders the model incomplete?
According to classical physics, the moving negatively charged electron should collapse into the positively charged nucleus. That is, it was not understood what kept the electron in its position away from the nucleus.
Describe briefly the main features of the Bohr model of the hydrogen atom.
Bohr proposed that the electron exists in discrete energy levels (shells), n, each with defined energies. These energy levels are said to be quantized, that is, the electron cannot exist “in between” the discrete energy levels. The electron exists predominantly in the lowest energy level (n=1), called “ground state”. If the atom is provided with an appropriate amount of energy, the electron will be “excited” to a higher energy level. An electron in higher energy level is unstable, and the electron will drop to lower energy levels and ultimately to ground state. As the electron returns to ground state, energy is lost by the atom as a photon, with an energy that exactly matches the energy difference between the higher and lower energy levels.
Define ground state.
In the Bohr model this is the lowest and most stable energy state for the electron. The electron exists predominantly in this state.
Define excited state.
An atom is in an excited state when it has absorbed energy which has allowed the electron to exist temporarily in one of the permitted higher energy states.
Define emission spectrum.
Discrete lines at designated wavelengths (in the UV, visible and/or IR) which result from electrons “falling” from higher to lower energy levels. Each element will have its own emission spectrum, which a reflection of the energy difference between excited states and ground state. By contrast, a continuous spectrum shows a continuous set of “lines” at all wavelengths.
Define absorption spectrum.
Discrete blanks (black lines) in a continuous spectrum at designated wavelengths (in the UV, visible and/or IR) which result from electrons “jumping” from lower to higher energy levels. Each element will have its own absorption spectrum, which a reflection of the energy difference between excited states and ground state. Absorption spectra are useful in identifying elements in stars.
What experimental evidence supported the concept of quantized energy states in the hydrogen atom?
An important piece of experimental evidence is the hydrogen emission spectrum. For example four discrete lines are easily discerned in the visible portion of the hydrogen emission spectrum. The wavelength of these lines correspond to energies which exactly match the expected energy difference (as calculated by Bohr) if an electron was “falling” from n=3 to n=2; n=4 to n=2; n=5 to n=2; n= 6 to n=2.
What are the strengths and weaknesses of the Bohr model?
The Bohr model introduces the important understanding that energy levels, where electrons are permitted to exist, are quantized. Unfortunately, experimental data and theoretical calculations based on the model do not work when applied to multi-electron atoms - that is, all other atoms besides hydrogen. Therefore, although it is assumed that quantized energy levels exist in all atoms, a more developed model is needed, especially to describe the electronic structure in multi-electron atoms. This more developed model is the quantum mechanical model.
Define orbital.
An orbital is a mathematical solution to the Schrodinger equation based on the assumption that electrons behave as waves. This mathematical solution is related to the 3D region of space outside the nucleus where there is a 90% probability of the electron existing. This mathematical solution is also related to the relative energy of the electron in the atom.
How many quantum numbers are needed to define an orbital? Describe each one.
Three quantum numbers are needed to define an orbital: (n) = principal quantum number- this is the most important quantum number in describing the energy state of the electron. This number also defines the shell; (l ) = azumithal quantum number - this describes shape of the region of space with high probability of finding the electron (i.e. the orbital). Essentially it describes the shape of the boundary surface diagram. This quantum is associated with the subshell designations s, p, d, f and therefore also defines, to some degree, the energy of the electron (s p d f).; (ml ) = magnetic quantum number - this describes the orientation in space of the orbital. For example, p subshell orbitals exist on the x, y, and z axes (i.e. perpendicular).
How many orbitals exist in the hydrogen atom?
There are potentially an infinite number of orbitals. According to the quantum model, the one electron in hydrogen exists primarily in the 1s orbital. This corresponds to ground state. This one electron however can be excited to higher energy orbitals (e.g. 2s, 3p, 4d etc).
How many quantum numbers are needed to define an electron?
Four quantum numbers are needed to describe a single electron. In addition to n, l, and ml, the fourth quantum number is called the spin quantum number (ms). The only allowed values for ms are +1/2 and -1/2. For Chem 112 students this is more easily understood as representing two possible electrons of exactly opposite spins. Electrons of different spins are often represented in orbital diagrams by arrows with the arrow tip either pointing upwards or downwards.
