Atomic structure

1. Atomic structure • To review, an atom consists of a small, dense nucleus containing all of its protons and neutrons, surrounded by electrons that fill the remaining volume of the atom. The atom stays electrically neutral because the number of protons and electrons are usually equal. In this section, the structure of atoms and the chemical properties will be studied, and the number of electrons and the way they are distributed generally determine them.

2. Atoms and Bonding Atoms of almost every element has the ability to combine with other atoms to form more complex structures. The forces of attraction that bind them together are chemical bonds.

3. Bonding • To understand chemistry, the nature and origin of chemical bonds is important, since the basis of chemical reactions is the forming and the breaking of bonds and the changes in bonding forces. There are two main classes of bonding forces: covalent bonds and ionic bonds. Covalent bonding deals with the sharing of electrons between atoms. Ionic bonding deals with the transfer of electrons between atoms.

4. Ionic Bonds • Ionic compounds (more commonly known as salts) are formed when metals react with nonmetals. A common salt, table salt, or sodium chloride (NaCl) is formed from a sodium ion, Na+ and a chloride ion, Cl-. A sodium atom, loses an electron and gives it to a chlorine atom. The sodium atom, is now an ion and has a positive charge, while the chlorine atom changes to an ion and has a negative charge. Once they are formed, the positive ion and negative ion attract each other in an ionic bond. The reason why electrons are transferred between these atoms is because for any stable compound to form from its elements, there must be a net lowering of potential energy (exothermic).

5. The Octet Rule • Basically, when ions are formed, they like to form a noble gas configuration. This is the octet rule, which states that atoms of most representative elements tend to gain or lose electrons until they have obtained a configuration that is the same as that of the nearest noble gas. All of the noble gases except helium have valence shells with eight electrons, which is why it is called the octet rule. So, the rule can be restated to say that atoms tend to gain or lose electrons until they have achieved an outer shell that contains an octet of electrons (eight electrons).

6. Exceptions to the Octet Rule • The octet rule works well for the representative metals (Group IA, IIA) and the nonmetals, but not for the transition elements and post-transiton elements. This is because they have d and f subshellorbitals. For example, tin (a post-transition metal) forms two ions, Sn2+ and Sn4+. The electron configurations are: Sn[Kr] 4d105s25p2 Sn2+ [Kr] 4d105s2 Sn4+ [Kr] 4d10 Neither of these configurations are noble gas configurations though. For iron ([Ar] 3d64s2), the 4s2 can be lost to make Fe2+. But Fe3+ can be made because a 3d5 (half-filled) subshell is more stable.

7. Ionic Lewis structures • Lewis symbols are used mostly for showing covalent bonds, but can be used to describe what happens during ion formation. For example, here is the formation of NaCl:

8. Covalent bonds • Most substances that we see are not ionic. They are not collections of electrically charged particles (ions), but are electrically neutral combinations called molecules. In ionic bonding, the potential energy lowering must occur. But many times, it is not possible because of the great range of IE. For example, lets take a hydrogen molecule (H2). As two H atoms approach each other, the electron of each atom feels the pull from the other atom. The electron density around each nucleus shifts to include both atoms. To lower the energy needed, the two nuclei share the electrons. Since the two electrons and the two nuclei repel each other but the electrons and nuclei attract each other, the two nuclei are kept a certain distance away from each other, called a bond length or bond distance. when the bond is formed, an amount of energy is release as the potential energies are lowered. This energy is called the bond energy.

9. Formation of a Hydrogen Molecule • More simply, an electron pair can be represented by a dash. • H-H Formulas drawn with Lewis symbols are called • Lewis formulas or Lewis structures, or more simply structural formulas. They show which atoms are present and how they are connected into each other

10. Covalent Bonding and the Octet Rule • Applied to covalent bonding, the octet rules states that when atoms form covalent bonds, they tend to share sufficient electrons so as to achieve an outer shell having eight electrons (except for hydrogen and helium, which have a stable outer shell of 2 electrons).

11. Covalent Bonding and the Octet Rule • For example, here is the bonding of hydrogen Chloride (HCl): • An octet is formed on the chlorine atom and the maximum pair on the hydrogen atom. Here is the bonding of a chlorine molecule (Cl2): • An octet is formed on the two chlorine atoms.

12. Covalent Bonding and the Octet Rule • Some atoms can also form more than one covalent bond. For example, carbon can form 4 bonds (because it has 4 valence electrons), nitrogen can form 3 (5 valence electrons), and oxygen can from 2 (6 valence electrons):

13. Double and Triple Bonds • So far, the bonds we've studied were single bondS, where only one pair of electrons were shared between atoms. To illustrate a triple bond, the nitrogen gas molecule will be used. A nitrogen atom looks like this, with 5 valence electrons around it. • When the two N atoms share electrons they have to share 3 pairs to add to their 5 electrons to complete their octet.