Plum pudding model
The atom is made up of positive, negative and neutral particles. Many theories were investigated to explain how these particles exist in the atom. It was agreed upon that the nucleus is located in the center, and it’s made up of the neutrons and protons.
One idea that was proposed is referred to as the plum pudding model. The positive charge (proton) is distributed all through the atom while the negative charges (electrons) were randomly existing in the atom.
Further investigation by other scientists, namely Niels Bohr and Ernest Rutherford, disproved the plum pudding model. Instead, research showed that the positive charge of the atom was existent only in the center at the nucleus. In addition, negative charges (electrons) orbited the nucleus.
The electrons exist in the atom in orbitals and in discrete, distinct energy levels. The atom has different energy levels for the electrons to occupy. The energy level is denoted by the letter n which can only be an integer starting at 1 and incrementing by 1. So, n can equal 1, 2, 3, etc.
In addition, there can be only a certain number of electrons at each energy level. The maximum number of electrons at each level is denoted by 2n2. So at the first energy level when n=1, there can be 2 (1)2 = 2 electrons. At the second energy level when n=2, there can be 2 (2)2 = 8 electrons. The lowest energy level, i.e., when n=1, is referred to as the ground state.
Electrons moving to higher states
Electrons can move from one energy level to another. For example, the hydrogen atom has one proton, one neutron and one electron. The electron is located in the ground state n=1. This electron can go to a higher state, implying it leaves the n=1 orbital and enters n=2 orbital.
When the electron accomplishes that task, it is then considered to be in an excited state. This state contains more energy. The electron can only achieve this migration by gaining energy to move to a higher state. This can take place when the electron receives a photon of light.
An electron can also drop to a lower energy state, thus releasing a photon. The photon energy will be equal to the energy lost by the atom. During the migration of the atom, the energy can be measured to determine the energy of the photon.
The first equation demonstrates that photon energy is equal to the energy in the excited state minus the energy in the ground state. The second equation shows another way to calculate the energy of the photon based on h = Planck’s constant, f = frequency, ʎ = wavelength and c = speed of light.
Ephoton = hƒ = hc / λ
The energy that is needed or released as electrons become excited or head toward the ground state, respectively, is not equal between the levels. From the ground state up, the space from one energy level to the next energy level decreases. So if an electron goes from level 4 to level 3, it will release less energy than if it were to go from level 3 to level 2.
Electron energy levels
The electron can go to higher and higher excited states. Each state has a higher and higher energy. This can be calculated by the following equation:
En = (–13.6/n2) eV
In this equation, e = electron charge, V = voltage, n = level and En = energy. Substituting this equation will assist in calculating the change in the energy of the electron as it traverses from one level to another. This will help determine the energy of the emitted photon.
The energy that is calculated is negative when referring to bound electrons. These are electrons that are closer to the nucleus, bound to the proton. As an electron goes to higher levels, the energy becomes less negative by increasing in value. The higher the level, the energy heads towards 0 since the electron is far from the nucleus, which would occur when n = ∞.
Emission and absorption spectrum
The electron that travels to a lower energy level will emit a photon which has a specific frequency and wavelength. Depending on the drop in the energy level, the photon that is emitted will have a different color. If the energy drop is small, not much energy will be released causing a photon of lower frequency to give off a red color. If the energy drop is large, much more energy will be released causing a photon of higher frequency to give off a blue/purple color.
This process creates a fingerprint for atoms. The emitted energy and light can be directed towards a prism which separates the light into bands of different colors creating a unique emission spectrum for each atom.
In the opposite manner, light can travel through an atom where electrons get excited and the exiting light can be refracted in the prism to see what colors are missing since those frequencies were absorbed by the electrons. This would create a unique absorption spectrum for each atom.
Atomic quantum numbers
The quantum numbers are used to describe the orbitals in which electrons can be found. The first one is n, which signifies the different energy levels for the electrons to occupy.
Each of these levels has different possible shapes. The shapes or orbitals are denoted by the letter l. The higher the level, the more the number of shapes. The number of shapes is limited per level by n-1. For example, when n = 2, l = 0 and 1. At n = 2, there are two possible shapes: circular or bilobed. Each of the shapes can have different possible orientations denoted by the letter ml. The orientations are also limited and range from –l to +l. The last quantum number is spin ms which can be only +1 or -1, up or down, respectively.
Pauli exclusion principle
The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. As long as one of the quantum numbers is different, the principle is being followed. For example, there are two electrons in an atom with the following numbers: 2, 0, 0 +1 and 2, 0, 0, -1. Both of these possibilities can exist since the spin number is different.
Electron structure notation
The electrons travel in different orbitals based on the l quantum number.
The shapes of the different l allow for creating a unique notation for each atom. If an atom is in the first energy level, n = 1, then l = 0 implies an s shape. This is denoted as 1s2 where 1 is the level, s is the shape, and 2 are the number of electrons that fill this level.
The next level that will be filled would be 2s2. s implies the spherical or circular orbital only. So this would be the level 2, the shape s and 2 electrons.
Now if the bilobed orbitals get filled at this level, then it would be referred to as 2p6. Here, the three bilobed orbitals would fill up, each carrying 2 electrons and giving a total of 6 electrons.
In this manner, d would have 5 different shapes (l = 2) each with 2 electrons giving d orbitals a maximum of 10 electrons. Finally, the f would have 7 different shapes (l = 3) each with 2 electrons giving f orbitals a maximum of 14 electrons.
Electron configuration table
The electron notation that is used allows all the atoms to be placed in a table called the electron configuration table or periodic table. Each part of the table represents the status of the different energy levels and orbitals.
For example, P for phosphorous. In order to write the electron notation, start at the top left and go to the right one line at a time. The notation would be 1s22s22p63s23p3. Since the P has only 3 electrons in the outer orbital, it can still acquire more electrons. The column to the far right is the noble gases column which are inert gases. These atoms have full outer orbitals. In addition, He is to the far right because the outer shell of s is full with 2 electrons. It is colored blue since it belongs to s orbital area.
Electron configuration and magnetism
Non-magnetic materials can respond to magnetic fields. If something is diamagnetic, electrons are usually paired inducing a repulsive force. If something is paramagnetic, there usually is an unpaired electron inducing an attractive force. As stated earlier, noble gases have full orbitals and paired electrons so they are referred to as diamagnetic.
Effective nuclear charge
The effective nuclear charge is the net positive charge that an electron experiences. Outer electrons see an effectively weakened positive charge due to the shielding from inner electrons at lower levels. The effective nuclear charge that the outer shell electron experiences is referred to as the core charge.
Each particle exists at a position with a particular momentum. Unfortunately, both of these measurements are not precise and rather uncertain. Heisenberg stated that the uncertainty principle helps mathematically estimate the physical properties of the particle in regards to position and momentum.
A particle’s position is described as the probability of location with a wave. The momentum is related to the wavelength of the uncertainty wave. However, to narrow down the location of the particle, construct many waves with varying wavelengths and momentum. Adding the waves together creates a new wave form from constructive interference and destructive interference. The new wave is called the wave packet and shows more of a centralized location of the position and momentum of the particle as seen in the figure below.
The top figure shows the uncertainty in momentum whereas the bottom figure shows the uncertainty in position. Both are inversely related. If one, uncertainty decreases; the other one will have an uncertainty that will increase.
The photoelectric effect
The photoelectric effect is the release or emission of electrons when light is shone onto an object, material or atom. Light can entirely release electrons from their orbitals. However, only high frequency light causes electrons to be released. When red light is used, no electrons are released. When higher energy blue light is used, electrons are able to be released.
The energy of release is called the work function (phi). The photon’s energy contributes partially to releasing the electron and partially to the kinetic energy of the released electron.