Electron Configuration
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1 Electron Configuration
2 Plumb Pudding Atomic Model Thomson s atomic model consisted of negatively charged electrons embedded in a ball of positive charge. Diagram pg 81 of chemistry text.
3 Rutherford s Model Electrons move in set paths around the nucleus much like planets around the sun. Rutherford s model could not explain why electrons are not pulled into the nucleus by the opposite charge. Rutherford s model was replaced two years later by a model proposed by Niels Bohr.
4 Bohr s Model of the Atom Electrons have a certain energy which determines the energy level of the electron Electrons must gain energy to move to a higher energy level or lose energy to move to a lower energy level. The difference in energy between two energy levels is known as a quantum of energy. Electrons can only be in an energy level, not in between energy levels Bohr s description of energy levels is still accurate today
5 Modern Model of the Atom The modern model explained that the electron behaves like both waves (energy) and like particles (matter). De Broglie explained that electrons could be waves confined to the space around the nucleus, and that the frequencies could correspond to the specific energy levels in which electrons are found.
6 Modern Model Cont. An electron s exact location, speed or direction cannot be determined because the electron is constantly moving and changing location. Electron clouds show the probable area in which an electron can be found. The darker the shading the more likely the electron is at that location
7 Waves Crest wavelength amplitude trough Frequency is the number of waves that pass a certain point per second Wavelength is the distance from one point on a wave to the exact same point On the next wave Amplitude is the distance from a crest or trough to the mid point of the wave As the frequency of a wave increases the wavelength decreases, therefore With an increase in energy there is an increase in frequency and decrease in wavelength.
8 Electromagnetic Spectrum All the wavelengths of electromagnetic radiation. Visible light spectrum are the wavelengths between 400nm (violet) and 700nm (red) The ultraviolet wavelengths are just below 400nm and the infrared wavelengths are just above 700nm.
9 Photoelectric Effect Explains the particle properties of electrons. This effect happens when light strikes a metal and electrons are released. The light must have a certain frequency in order to remove an electron from the metal. Einstein proposed that light is a stream of particles the energy of which determines the lights frequency. In order to remove the electron from the metal, light must have a minimum frequency.
10 Line-Emission Spectrum When a high-voltage current is passed through a tube of hydrogen gas at low pressure, lavender colored light is seen. When passed through a spectrum only a few wavelengths of light can be seen. This is called a line-emission spectrum. When electrons are excited by an energy source such as electricity or heat they move from a low energy level to a higher energy level.
11 Light Emission Cont. Upon reaching the higher energy level the electron then releases the absorbed energy in the form of electromagnetic radiation and moves back to the lower energy level. The lowest energy level occupied by an electron is called its ground state and the higher energy level is referred to as the excited state.
12 Light Emission cont. The energy that is released has a specific wavelength. Each move from one energy level to another releases light of a different wavelength. Elements can be identified by their lineemission spectrum The more energy released by the drop in energy levels the higher the frequency and the shorter the wavelength.
13 Quantum Numbers Principal Quantum Number tells the main energy level occupied by the electron The energy level is symbolized by the variable n. There are 7 energy levels therefore n=1-7. As n increases the electrons distance from the nucleus increases. The energy level closest to the nucleus has the lowest energy electrons.
14 Quantum Numbers Cont. Angular Momentum Quantum Number describes the shape of the electrons orbital. There are four shapes: s = spherical p = dumbbell shaped d = four lobed f = so complex it cannot be described
15 Orbital Quantum Numbers Cont. The 1 st energy level has only an s orbital in it The 2 nd energy level has s and p orbitals The 3 rd energy level has s, p and d orbitals The 4 th through 7 th energy levels have s, p, d and f orbitals. The number of orbitals in an energy level can be calculated by the square of the energy level ie. 1 2 = 1 orbital in 1 st energy level, 2 2 = 4 orbitals in the 2 nd energy level, s, p x, p y, p z
16 Magnetic Quantum Numbers The Mangentic Quantum Numbers describe the orientation in space of the orbitals on the x, y and z axis. The s orbital has only one orientation in each energy level encompassing all the planes with one orbital. The p orbital has 3 orientations in each energy level above the 1 st energy level. The d orbital has 5 orientations in each energy level above the 2 nd energy level. The f orbital has 7 orientations in each energy level above the 3 rd energy level.
17 Spin Quantum Number The Spin Quantum Number describes the orientation of the electron s magnetic field relative to an outside magnetic field It is symbolized by +1/2 or -1/2. It can also be described as the spin of the electron being either clockwise or counter clockwise. A single orbital can hold a maximum of 2 electrons, which must have opposite spins
18 Electrons Electrons are found in energy levels The number of electrons in each energy level: 2n 2 = the number of electrons per energy level 1 st 2, 2 nd 8, 3 rd 18, 4 th 7 th 32 each Orbitals are regions within energy levels where electrons are found, they are sometimes called electron clouds. A maximum 2 electrons can occupy each orbital Electron energy increases moving from 1 st through 7 th energy levels and s-f orbitals
19 Pauli Exclusion Principle No two electrons of the same atom can have the same exact set of quantum numbers. They can have the same values except for the spin quantum number it has to be different.
20 Hund s Rule Orbitals of the same energy are occupied by one electron until all orbitals have an electron then a second electron can be placed in the orbital. Example: In the 2 p orbitals for the element Nitrogen there is only one electron in each of the 3 orbitals.
21 Aufbau Principle German for building up Electrons fill orbitals beginning with the lowest energy first. Use the Orbital Filling handout by following the arrows from tail to head to show the order in which electrons fill orbitals
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