Arrangement of Electrons in Atoms

Gap-Fill Exercise
Content © 2006 Melanie Cecere; Authors of Activity: Melanie Cecere & Tami Maloney; All rights reserved. No commercial, for-profit use of this material is allowed. E-mail comments and questions to Tami Maloney.

Fill in all the gaps, then press "Check" to check your answers. Use the "Hint" button to get a free letter if an answer is giving you trouble. You can also click on the "[?]" button to get a clue. Note that you will lose points if you ask for hints or clues!
Development of a New Atomic Model. The Rutherford model of the atom did not explain why the electron wasn't drawn into the of the atom. Early in the 20th century, a new model evolved as a result understanding the relationship between light and an atom's electrons.
A. Properties of light. Before 1900, scientists thought of light as acting as only. New information led them to realize that light has characteristics of , as well.

1. Light as Waves. Electromagnetic radiation (EMR) is a form of that exhibits wavelike behavior as it travels through space. The electromagnetic spectrum encompasses all the forms of electromagnetic radiation. is the distance between corresponding points on adjacent waves. is the number of waves that pass a given point in a specific time, usually one second. It is measured in (Hz). As the wavelength of light decreases, its frequency , and vice versa.
2. Light as particles. The Effect is the emission of electrons from a metal when light shines on the metal. Electromagnetic radiation strikes the surface of the metal, ejecting from the metal and creating an electric current. In order for an electron to be ejected from a metal surface, the electron must be struck by a single possessing at least the minimum energy required to knock the electron loose. The theory of light suggested that ANY frequency of light could eject an electron, but the photoelectric effect clearly shows this is wrong. German physicist Max proposed and explanation. He suggested light acts as particles-energy is emitted in small, specific amounts called . A quantum is the minimum quantity of energy that can be lost or gained by an .

B. The Hydrogen-Atom Line-Emission Spectrum. The state is the lowest energy state of an atom. The state is the state in which an atom has a higher potential energy than in its ground state. When an excited atom returns to ground state, it gives off the energy it gained in its excited state in the form of radiation. When electric current is passed through a vacuum tube filled with hydrogen, a pink light is produced. When the pink light passes through a prism, it produced specific, unique pattern of wavelengths of light known as hydrogen's line-emission . The greater the energy jumps, the shorter the . Attempts to explain these observations led to the theory of the atom.

C. Bohr Model of the Hydrogen Atom. Niels was a Danish physicist. He proposed a model of the hydrogen atom that linked the atom's electron with emissions. In his model, the electron can circle the nucleus only in certain paths, or orbits. Energy is involved to change paths and photons are emitted if the path drops. This model accounts mathematically for the energy of each transition in the hydrogen line-emission spectrum. BUT, Bohr's model of quantitized energy levels could not predict the energy levels of electrons in atoms with more than 1 .

The Quantum Model of the Atom
A. Electrons as Waves - French scientist Louis said:
- Electrons behave as waves and can only exist at specific . These frequencies correspond to specific energies -- the quantized energies of Bohr's orbits. In 1924, de Broglie said if waves can have particle-like character, maybe particles can have character.
- His ideas were proven in 1927 when investigators demonstrated that electrons can be bent or defracted, like light (defraction), and that electron beams can interfere or overlap with each other like light (interference). This overlapping causes a reduction of energy in some areas and increase in others.
B. The Uncertainty Principle. German theoretical physicist Werner Heisenberg proposed this principle in 1927. Electrons can be detected by their interaction with photons but, because they each have about the same energy, the photon causes the to be knocked off course. As a result there is always a basic in trying to locate and electron. The principle states that it is to determine simultaneously the position and velocity of an electron or any other particle. The reason is that in order to observe a particle, you must interact with it or illuminate it. In order to "see" it, you would need to use photons of light. The photons would change the velocity of the electron. Any attempt to observe the location and speed fails because either its location or speed changes during the attempt. This was difficult for the scientific community to accept, but it has proven to be one of the fundamental principles of our present understanding of light and matter.
C. The Schrodinger Equation. Austrian physicist Erwin Schrodinger used the hypothesis that electrons have a dual wave-particle nature to develop an equation that treated electrons in atoms as waves. Only electrons of specific frequencies () solved the equation. This provided yet another foundation for the theory. The theory describes mathematically the wave properties of electrons and other very small particles. It holds that electrons do not travel around the nucleus in neat orbits as Bohr stated, but they exist in regions called . An orbital is a three-dimensional region around the nucleus that indicates the location of an electron.
- Probability and Energy Levels: The wave equations of quantum mechanics allow scientists to determine the probability of finding a particle at a particular place at a particular time. The equations do not calculate exact orbits but, instead, regions of space inside the atoms where an is likely to be at any time (orbital). If you plot all the probabilities, you will get diffuse with regions of high and low density. The theory of wave mechanics states that a wave representing the electron must "fit" inside the atom in such a way so that it meets itself without any overlap (standing wave). The wavelength depends on the energy of the . There are only certain energies for which the wavelength's are just right to form standing waves in the atom. These energies correspond to energy levels in Bohr's atom.
Wave mechanics gives the of finding an electron, not an orbit or path like Bohr thought. The average position of the plot of probability points is a spherical shell centered on the nucleus ( levels). The energy levels or shells are called energy levels. The number of the shell or energy level is called the principle quantum level (n).
Energy levels have sublevels. The energy of each sublevel within a level is slightly different. The existence of accounts for the abundance of lines in the spectra of atoms. (Bohr's model couldn't explain them because he only had principal energy levels). The number of sublevels in any energy level is the same as its principle number.
n = 1 sublevel 1s
n = 2 sublevels 2s 2p
n = 3 sublevels 3s 3p 3d
n = 4 sublevels 4s 4p 4d 4f
The lowest sublevel in each principal level is called the sublevel. The next lowest sublevel is called the sublevel; the next lowest is the sublevel and the highest is the sublevel. Any higher (g, h, i, etc.) are not needed yet-not enough elements.
are regions within a sublevel or an energy level where electrons may be found. There is a maximum of electrons per orbital.
s sublevel has: 1 orbital = electrons max
p sublevel has: 3 orbitals = electrons max
d sublevel has: orbitals = 10 electrons max
f sublevel has: orbitals = 14 electrons max

D. Atomic Orbitals and Quantum Numbers. Quantum specify the properties of atomic orbitals and the properties of electrons in orbitals. A set of numbers is used to describe the location of an electron (its address) - n, l, m, s
- The principle quantum number is symbolized by , and indicates the main energy level occupied by the electron. Values are positive integers. As n increases, the energy and from the nucleus increases. More than one electron can have the same n. Theses electrons are in the same shell and the total number of orbitals that exist in a given shell or main energy level is n2. Electrons cannot have energies corresponding to anything by values of n.
- The momentum quantum number, symbolized by , indicates the of the orbital or the sublevel. Except for the 1st energy level, orbitals of different shapes (sublevels) exist for a given n. For a specific main energy level, the number of shapes possible is equal to n. The values of l allowed are zero and all possible integers less than or equal to n-1. An orbital is assigned a letter:
orbitals are spherical
orbitals are dumbbell shaped
d orbitals are more complex
f orbitals are really complex
If n = 1, there is one sublevel, the s orbital. If n = 2, there are sublevels, the s and p orbitals. If n = 3, there are sublevels.
Each atomic orbital is designated by the principle quantum number followed by the letter of the sublevel. "1s" is the s orbital in the main energy level. "2p" is the p sublevel in the main energy level. "4d" designates the sublevel in the 4th main energy level.
- The quantum number, symbolized by m, indicates the orientation of an orbital around the nucleus. For example, px, py, and pz all designate the p orbital, but on each of 3 different axes. There is only possible orientation for the s orbital (m = 0). There are 3 for the orbital (m = -1, 0, or +1); for the d orbital (m = -2, -1, 0, +1, or +2) and 7 for the f orbital (m = -3, -2, -1, 0, +1, +2, or +3).
- The quantum number has only two possible values (+ ½ or - ½ ) which indicate the two fundamental spins states of an electron in an orbital. A single orbital can hold a maximum of electrons, which must have opposite spins.

The electron configuration is the arrangement of electrons in an atom. Ground-state electron configuration is the energy arrangement of the electrons for each element.
A. Rules for Electron Configuration: To build up electron configurations for the ground state of any particular atom, first the energy levels are determined, then electrons are added to the orbitals one by one according to three basic rules:
1. principle: an electron occupies the lowest-energy orbital that can receive it.
2. exclusion principle: no two electrons in the same atom can have the same set of four quantum numbers (importance of spin quantum numbers).
3. rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin (placing unpaired electrons in as many separate orbitals as possible in the same sublevels). Analogy: strangers on a bus sit in separate seats until they have to double up.

B. Order of filling: The sublevels do not full up in numerical order. There is a more configuration at a half-filled or filled sublevel. Using the Periodic Table of the Elements, the order in which shells and subshells are filled can be seen by following the table from left to right across each period. The elements are arranged in the same way orbitals are filled: 1st, the 1s shell gets 2 electrons, then the 2s subshell holds two more, the 2p subshell has 3 for a total of 7 electrons. After the third period, the filling of subshells becomes more complicated. For example, the 4s subshell has a lower energy level that the 3d subshell. As orbitals are filled, the 4s comes before 3d and then the 4p is filled last. This is easy to remember because the atomic numbers increase in the same order as the subshells are filled: 4s, 3d, 4p.
C. Notation: There are three basic ways to note electron configuration:

1. Notation: Unoccupied orbitals are represented by a line with the principle quantum number and sublevel written underneath.
2. Electron-configuration Notation. No lines, no arrows. The number of electrons is a is shown by adding a superscript to the sublevel designation.
3. Notation. This is an abbreviated system. First, you find the element on the Periodic Chart. Then go to the row above and put the atomic symbol for the element in the last row-far right (the noble gas of the period) in brackets. Then start writing your electron configuration for that element.