Atoms are the basic units of matter and the defining structure of elements. Atoms are made up of three particles: protons, neutrons and electrons.
Protons and neutrons are heavier than electrons and reside in the center of the atom, which is called the nucleus. Electrons are extremely lightweight and exist in a cloud orbiting the nucleus. The electron cloud has a radius 10,000 times greater than the nucleus.
Protons and neutrons have approximately the same mass. However, one proton weighs more than 1,800 electrons. Atoms always have an equal number of protons and electrons, and the number of protons and neutrons is usually the same as well. Adding a proton to an atom makes a new element, while adding a neutron makes an isotope, or heavier version, of that atom.
Nucleus
The nucleus was discovered in 1911, but its parts were not identified until 1932. Virtually all the mass of the atom resides in the nucleus. The nucleus is held together by the “strong force,” one of the four basic forces in nature. This force between the protons and neutrons overcomes the repulsive electrical force that would, according to the rules of electricity, push the protons apart otherwise.
Protons
Protons are positively charged particles found within atomic nuclei. They were discovered by Ernest Rutherford in experiments conducted between 1911 and 1919.
The number of protons in an atom defines what element it is. For example, carbon atoms have six protons, hydrogen atoms have one and oxygen atoms have eight. The number of protons in an atom is referred to as the atomic number of that element. The number of protons in an atom also determines the chemical behavior of the element. The Periodic Table of the Elements arranges elements in order of increasing atomic number.
Protons are made of other particles called quarks. There are three quarks in each proton — two “up” quarks and one “down” quark — and they are held together by other particles called gluons.
Electrons
Electrons have a negative charge and are electrically attracted to the positively charged protons. Electrons surround the atomic nucleus in pathways called orbitals. The inner orbitals surrounding the atom are spherical but the outer orbitals are much more complicated.
An atom’s electron configuration is the orbital description of the locations of the electrons in an unexcited atom. Using the electron configuration and principles of physics, chemists can predict an atom’s properties, such as stability, boiling point and conductivity.
Typically, only the outermost electron shells matter in chemistry. The inner electron shell notation is often truncated by replacing the long-hand orbital description with the symbol for a noble gas in brackets. This method of notation vastly simplifies the description for large molecules.
For example, the electron configuration for beryllium (Be) is 1s22s2, but it’s is written [He]2s2. [He] is equivalent to all the electron orbitals in a helium atom. The Letters, s, p, d, and f designate the shape of the orbitals and the superscript gives the number of electrons in that orbital.
Neutrons
Neutrons are uncharged particles found within atomic nuclei. A neutron’s mass is slightly larger than that of a proton. Like protons, neutrons are also made of quarks — one “up” quark and two “down” quarks. Neutrons were discovered by James Chadwick in 1932.
RULES FOR FILLING ORBITALS
Rule 1 – Lowest energy orbitals fill first. Thus, the filling pattern is 1s, 2s, 2p, 3s, 3p, 4s, 3d, etc. Since the orbitals within a subshell are degenerate (of equal energy), the entire subshell of a particular orbital type is filled before moving to the next subshell of higher energy.
Rule 2 – Pauli Exclusion Principle – Only two electrons are permitted per orbital and they must be of opposite spin. If one electron within an orbital possesses a clockwise spin, then the second electron within that orbital will possess a counterclockwise spin. Two electrons with opposite spins found in the same orbital are referred to as being paired.
Rule 3- Hund’s Rule – The most stable arrangement of electrons in a subshell occurs when the maximum number of unpaired electrons exist, all possessing the same spin direction. This occurs due to the degeneracy of the orbitals, all orbitals within a subshell are of equal energy. Electrons are repulsive to one another and only pair after all of the orbitals have been singly filled.
Rules for Assigning Electron Orbitals
Occupation of Orbitals
Electrons fill orbitals in a way to minimize the energy of the atom. Therefore, the electrons in an atom fill the principal energy levels in order of increasing energy (the electrons are getting farther from the nucleus). The order of levels filled looks like this:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p
One way to remember this pattern, probably the easiest, is to refer to the periodic table and remember where each orbital block falls to logically deduce this pattern. Another way is to make a table like the one below and use vertical lines to determine which subshells correspond with each other.
Pauli Exclusion Principle
The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. The first three (n, l, and ml) may be the same, but the fourth quantum number must be different. A single orbital can hold a maximum of two electrons, which must have opposing spins; otherwise they would have the same four quantum numbers, which is forbidden. One electron is spin up (ms = +1/2) and the other would spin down (ms = -1/2). This tells us that each subshell has double the electrons per orbital. The s subshell has 1 orbital that can hold up to 2 electrons, the p subshell has 3 orbitals that can hold up to 6 electrons, the d subshell has 5 orbitals that hold up to 10 electrons, and the f subshell has 7 orbitals with 14 electrons.
Example 1: Hydrogen and Helium
The first three quantum numbers of an electron are n=1, l=0, ml=0. Only two electrons can correspond to these, which would be either ms = -1/2 or ms = +1/2. As we already know from our studies of quantum numbers and electron orbitals, we can conclude that these four quantum numbers refer to the 1s subshell. If only one of the ms values are given then we would have 1s1 (denoting hydrogen) if both are given we would have 1s2 (denoting helium). Visually, this is be represented as:
Hund’s Rule
When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy (also referred to as degenerate) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. When visualizing this processes, think about how electrons are exhibiting the same behavior as the same poles on a magnet would if they came into contact; as the negatively charged electrons fill orbitals they first try to get as far as possible from each other before having to pair up.
The Aufbau Process
Aufbau comes from the German word “aufbauen” meaning “to build.” When writing electron configurations, orbitals are built up from atom to atom. When writing the electron configuration for an atom, orbitals are filled in order of increasing atomic number. However, there are some exceptions to this rule.
Example 3: 3rd row elements
Following the pattern across a period from B (Z=5) to Ne (Z=10), the number of electrons increases and the subshells are filled. This example focuses on the p subshell, which fills from boron to neon.
• B (Z=5) configuration: 1s2 2s2 2p1
• C (Z=6) configuration:1s2 2s2 2p2
• N (Z=7) configuration:1s2 2s2 2p3
• O (Z=8) configuration:1s2 2s2 2p4
• F (Z=9) configuration:1s2 2s2 2p5
• Ne (Z=10) configuration:1s2 2s2 2p6
Exceptions
Although the Aufbau rule accurately predicts the electron configuration of most elements, there are notable exceptions among the transition metals and heavier elements. The reason these exceptions occur is that some elements are more stable with fewer electrons in some subshells and more electrons in others (Table 1).
Table 1: Exceptions to Electron Configuration Trends
Period 4: Period 5:
Chromium: Z:24 [Ar] 3d54s1
Niobium: Z:41 [Kr] 5s1 4d4
Copper: Z:29 [Ar] 3d104s1
Molybdenum: Z:42 [Kr] 5s1 4d5
Ruthenium: Z:44 [Kr] 5s1 4d7
Rhodium: Z:45 [Kr] 5s1 4d8
Palladium: Z:46 [Kr] 4d10
Silver: Z:47 [Kr] 5s1 4d10
Period 6: Period 7:
Lanthanum: Z:57 [Xe] 6s2 5d1
Actinium: Z:89 [Rn] 7s2 6d1
Cerium: Z:58 [Xe] 6s2 4f1 5d1 Thorium: Z:90 [Rn] 7s2 6d2
Gadolinium: Z:64 [Xe] 6s2 4f7 5d1 Protactium: Z:91 [Rn] 7s2 5f2 6d1
Platinum: Z:78 [Xe] 6s1 4f14 5d9 Uranium: Z:92 [Rn] 7s2 5f3 6d1
Gold: Z:79 [Xe] 6s1 4f14 5d10
Neptunium: Z:93 [Rn] 7s2 5f4 6d1
Curium: Z:96 [Rn] 7s2 5f7 6d1
Lawrencium: Z:103 [Rn] 7s2 5f14 7p1
Writing Electron Configurations
When writing an electron configuration, first write the energy level (the period), then the subshell to be filled and the superscript, which is the number of electrons in that subshell. The total number of electrons is the atomic number, Z. The rules above allow one to write the electron configurations for all the elements in the periodic table.
Three methods are used to write electron configurations:
1. orbital diagrams
2. spdf notation
3. noble gas notation
Each method has its own purpose and each has its own drawbacks.
Orbital Diagrams
An orbital diagram, like those shown above, is a visual way to reconstruct the electron configuration by showing each of the separate orbitals and the spins on the electrons. This is done by first determining the subshell (s,p,d, or f) then drawing in each electron according to the stated rules above.
spdf Notation
The most common way to describe electron configurations is to write distributions in the spdf notation. Although the distributions of electrons in each orbital are not as apparent as in the diagram, the total number of electrons in each energy level is described by a superscript that follows the relating energy level. To write the electron configuration of an atom, identify the energy level of interest and write the number of electrons in the energy level as its superscript as follows: 1s2. This is the electron configuration of helium; it denotes a full s orbital. The periodic table is used as a reference to accurately write the electron configurations of all atoms.
Noble Gas Notation
This brings up an interesting point about elements and electron configurations. As the p subshell is filled in the above example about the Aufbau principle (the trend from boron to neon), it reaches the group commonly known as the noble gases. The noble gases have the most stable electron configurations, and are known for being relatively inert. All noble gases have their subshells filled and can be used them as a shorthand way of writing electron configurations for subsequent atoms. This method of writing configurations is called the noble gas notation, in which the noble gas in the period above the element that is being analyzed is used to denote the subshells that element has filled and after which the valence electrons (electrons filling orbitals in the outer most shells) are written. This looks slightly different from spdf notation, as the reference noble gas must be indicated.
Example 6: Vanadium
What is the electronic configuration of vanadium (V, Z=23)?
SOLUTION
Vanadium is the transition metal in the fourth period and the fifth group. The noble gas preceding it is argon (Ar, Z=18), and knowing that vanadium has filled those orbitals before it, argon is used as the reference noble gas. The noble gas in the configuration is denoted E, in brackets: [E]. To find the valance electrons that follow, subtract the atomic numbers: 23 – 18 = 5. Instead of 23 electrons to distribute in orbitals, there are 5. Now there is enough information to write the electron configuration:
Vanadium, V: [Ar] 4s2 3d3
This method streamlines the process of distributing electrons by showing the valence electrons, which determine the chemical properties of atoms. In addition, when determining the number of unpaired electrons in an atom, this method allows quick visualization of the configurations of the valance electrons. In the example above, there are a full s orbital and three half filled d orbitals.
EVALUATION
1.Write the electronic configuration of the first twenty elements.
2.Write out the characteristics of the first three fundamental particles in an atom.
3.An atom of an element is represented by X. How many electrons,protons and neutrons are in the atom? Write the electronic structure of the atom.
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