2p orbitals. These three 2p orbitals are designated as 2px, 2py and 2pz. The third shell contains one 3s orbital, three 3p orbitals and five 3d orbitals. Thus, the first shell can hold only two 18 CH2 ATOMIC STRUCTURE AND BONDING electrons, the second shell eight electrons and the third shell up to 18 electrons, and so on. As the number of electrons goes up, the shell numbers also increase. Therefore, electron shells are identified by the principal quantum number, n¼ 1, 2, 3 and so on. The electronic configuration of an atom describes the number of electrons that an atom possesses, and the orbitals in which these electrons are placed. The arrangements of electrons in orbitals, subshells and shells are called electronic configurations. Electronic configurations can be repre- sented by using noble gas symbols to show some of the inner electrons, or by using Lewis structures in which the valence electrons are represented by dots. Valence is the number of electrons an atom must lose or gain to attain the nearest noble gas or inert gas electronic configuration. Electrons in the outer shells that are not filled are called valence electrons. The ground-state electronic configuration is the lowest energy, and the excited-state electronic configuration is the highest energy orbital. If energy is applied to an atom in the ground state, one or more electrons can jump into a higher energy orbital. Thus, it takes a greater energy to remove an electron from the first shell of an atom than from any other shells. For example, the sodium atom has electronic configuration of two, eight and one. Therefore, to attain the stable configuration, the Na atom must lose one electron from its outermost shell and become the nearest noble gas configuration, i.e. the configuration of neon, which has the electronic configuration of two and eight. Thus, sodium has a valence of 1. Since all other elements of Group IA in the periodic table have one electron in their outer shells, it can be said that Group IA elements have a valence of 1. At the far end on the right hand side of the periodic table, let us take another example, chlorine, which has the electronic configuration of two, eight and seven, and the nearest noble gas is argon, which has the electronic configuration of two, eight and eight. To attain the argon electronic configuration chlorine must gain one electron. Therefore, chlorine has a valence of 1. Since all other elements of Group 7A in the periodic table have seven electrons in their outermost shells and they can gain one electron, we can say that the Group 7A elements have a valence of 1. Shell Total number of shell electrons Relative energies of shell electrons 4 32 Higher 3 16 \ufffd\ufffd \ufffd\ufffd!2 8 1 2 Lower 2.2 ATOMIC STRUCTURE: ORBITALS AND ELECTRONIC CONFIGURATIONS 19 Each atom has an infinite number of possible electronic configurations. We are here only concerned with the ground-state electronic configuration, which has the lowest energy. The ground-state electronic configuration of an atom can be determined by the following three principles. \ufffd The Aufbau principle states that the orbitals fill in order of increasing energy, from lowest to highest. Because a 1s orbital is closer to the nucleus it is lower in energy than a 2s orbital, which is lower in energy than a 3s orbital. \ufffd The Pauli exclusion principle states that no more than two electrons can occupy each orbital, and if two electrons are present, their spins must be paired. For example, the two electrons of a helium atom must occupy the 1s orbital in opposite spins. \ufffd Hund\u2019s rule explains that when degenerate orbitals (orbitals that have same energy) are present but not enough electrons are available to fill all the shell completely, then a single electron will occupy an empty orbital first before it will pair up with another electron. This is understandable, as it takes energy to pair up electrons. Therefore, the six electrons in the carbon atom are filled as follows: the first four electrons will go to the 1s and 2s orbitals, a fifth electron goes to the 2px, the sixth electron to the 2py orbital and the 2pz orbital will remain empty. The ground-state electronic configurations for elements 1\u201318 are listed below (electrons are listed by symbol, atomic number and ground-state electronic configuration). Shell Number of orbitals contained each shell 4 4s, 4px, 4py, 4pz, five 4d, seven 4f 3 3s, 3px, 3py, 3pz, five 3d 2 2s, 2px, 2py, 2pz 1 1s First period Second period Third period H 1 1s1 Li 3 [He] 2s1 Na 11 [Ne] 3s1 He 2 1s2 Be 4 [He] 2s2 Mg 12 [Ne] 3s2 B 5 [He] 2s2 2p1 Al 13 [Ne] 3s2 3p1 C 6 [He] 2s2 2p2 Si 14 [Ne] 3s2 3p2 7 [He] 2s2 2p3 P 15 [Ne] 3s2 3p3 8 [He] 2s2 2p4 S 16 [Ne] 3s2 3p4 9 [He] 2s2 2p5 Cl 17 [Ne] 3s2 3p5 10 [He] 2s2 2p6 Ar 18 [Ne] 3s2 3p6 20 CH2 ATOMIC STRUCTURE AND BONDING Let us see how we can write the ground-state electronic configurations for oxygen, chlorine, nitrogen, sulphur and carbon showing the occupancy of each p orbital. Oxygen has the atomic number 8, and the ground-state electronic configuration for oxygen can be written as 1s2 2s2 2px 2 2py 1 2pz 1. Similarly, we can write the others as follows: Chlorine (atomic number 17): 1s2 2s2 2px 2 2py 2 2pz 2 3s2 3px 2 3py 2 3pz 1 Nitrogen (atomic number 7): 1s2 2s2 2px 1 2py 1 2pz 1 Sulphur (atomic number 16): 1s2 2s2 2px 2 2py 2 2pz 2 3s2 3px 2 3py 1 3pz 1 Carbon (atomic number 6): 1s2 2s2 2px 1 2py 1 2pz 0 2.3 Chemical bonding theories: formation of chemical bonds Atoms form bonds in order to obtain a stable electronic configuration, i.e. the electronic configuration of the nearest noble gas. All noble gases are inert, because their atoms have a stable electronic configuration in which they have eight electrons in the outer shell except helium (two electrons). Therefore, they cannot donate or gain electrons. One of the driving forces behind the bonding in an atom is to obtain a stable valence electron configuration. A filled shell is also known as a noble gas configuration. Electrons in filled shells are called core electrons. The core electrons do not participate in chemical bonding. Electrons in shells that are not completely filled are called valence electrons, also known as outer-shell electrons, and the energy level in which they are found is also known as the valence shell. Carbon, for example, with the ground-state electronic configuration 1s2 2s2 2p2, has four outer-shell electrons. We generally use the Lewis structure to represent the outermost electrons of an atom. 2.3.1 Lewis structures Lewis structures provide information about what atoms are bonded to each other, and the total electron pairs involved. According to the Lewis theory, an atom will give up, accept or share electrons in order to achieve a filled outer shell that contains eight electrons. The Lewis structure of a covalent molecule shows all the electrons in the valence shell of each atom; the bonds between atoms are shown as shared pairs of electrons. Atoms are most 2.3 CHEMICAL BONDING THEORIES: FORMATION OF CHEMICAL BONDS 21 stable if they have a filled valence shell of electrons. Atoms transfer or share electrons in such a way that they can attain a filled shell of electrons. This stable configuration of electrons is called an octet. Except for hydrogen and helium, a filled valence shell contains eight electrons. Lewis structures help us to track the valence electrons and predict the types of bond. The number of valence electrons present in each of the elements is to be considered first. The number of valence electrons determines the number of electrons needed to complete the octet of eight electrons.