The interaction between atoms that leads to a rearrangement of the electron to a more stable state is what we define as a chemical bond.
All atoms, except those of the noble gases*, readily engage in chemical bonding either with atoms of their own kind (in elements) or with atoms of a different kind (in compounds).
3 Types of chemical bonding:
- Ionic bond - electrostatic forces hold together oppositely charged ions.
- Covalent bond - two atoms sharing electron pair(s). Electrons are mutually attracted by the nuclei of adjacent atoms.
- Metallic bond -is a complicated type of bond. At the simples level metal atoms are envisaged of being metal cations in a “sea” of electrons. The electrons are shared between all the ions simultaneously. A better description of metals is obtained from the band theory.
Ionic bond:
Ionic bonds form between oppositely charged ions, generally, but not always, between metals and non-metals.
Metals versus Nonmetals:
Metals are conductors of heat and electricity, have a lustre (shiny) appearance, are ductile (can be drawn in to wirers) and are malleable (can be hammered out into thin foils). With the exception of hydrogen, which is a non-metal, metals are located to the left in the periodic table.
Non-metals are generally non conductors, do not have a metallic lustre and are brittle. They are located to the right in the periodic table.
The elements near the “step ladder” have both, metallic and non-metallic properties and are called metalloids or semi metals.
To understand chemical bonding trends in three atomic properties, we must be considered: electron affinity, ionization energy and atomic radius
Trends in atomic radii can be rationalized if we consider the shielding or screening effect of the core electrons.
The more electronic shells in an atom, the larger the atom. Atomic radius increases from top to bottom in a group.
The atomic radius decreases from left to right through a period of elements.
Ionic Radius:
Periodic Trends in Atomic Radii
Trends in atomic radii can be rationalized if we consider the shielding or screening effect of the core electrons.
Effective nuclear charge:
Zeff = Z – S
Z = atomic number, S = number of inner electrons that screen an outer electronThe more electronic shells in an atom, the larger the atom. Atomic radius increases from top to bottom in a group.
The atomic radius decreases from left to right through a period of elements.
Ions
Ions are charged atoms or groups of atoms.Ionic Radius:
- Cations are smaller than the atoms from which they are formed. Isoelectronic cations have the same number of electrons in identical configurations. Na+ and Mg2+ both have closed shells (both have 10 electrons but the magnesium nucleus has more protons to attract the 10 electrons then the sodium nucleus). For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius.
- Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the charge, the larger the ionic radius.
Electron affinity and ionization energy:
Electron Affinity: As we have seen an atom can gain one electron or more. This fundamental atomic property is known as electron affinity and is usually exothermic.
Nonmetals which are located to the right in the periodic table do not lose but tend to gain electrons and form anions. We called this atomic property electron affinity and the following trend is observed within the periodic table: Electron affinity generally increases from the left to the right within a period and decreases from the top the bottom in a group.
Ionization Energy: An atom can also lose one electron (or more). This fundamental atomic property is called ionization or ionization energy. Atoms are stable species. It requires energy to release an electron from an atom so ionization is an endothermic process and ionization energies are positive
Metals tend to lose electrons and form cations
I1 is called the first ionization energy, I2 the 2nd ionization energy.
Trends of ionization energy within a group in the periodic table: As you go down a group (chemical family) the ionization energy decreases. It is not difficult to understand the rationale behind that. With every new row (period) in the periodic table a new electronic shell is being filled with electrons (atoms get bigger!) The further away the electron is from the nucleus, the smaller is its attractive interaction with the nucleus and the easier it is to remove it.
Trends of ionization energy within a period: As you go from the left to the right ionization energy increases. Electrons enter the same shell within a perioid. This leads to an increase in repulsion between the electrons in the outermost electronic shell. At the same time as the number of electrons increases so does the number of protons in the nucleus and the greater the positive charge of the nucleus the greater the attraction towards the electrons. It is this increased attraction that dominates over electron repulsion. Thus ionization energies generally increase from the left to the right within a period. We will take a closer look at the exceptions later. This nucleus having a “tighter grip” on the outermost electrons also means that atomic radius decreases from the left to the right in a period.
Trends of ionization energy within a period: As you go from the left to the right ionization energy increases. Electrons enter the same shell within a perioid. This leads to an increase in repulsion between the electrons in the outermost electronic shell. At the same time as the number of electrons increases so does the number of protons in the nucleus and the greater the positive charge of the nucleus the greater the attraction towards the electrons. It is this increased attraction that dominates over electron repulsion. Thus ionization energies generally increase from the left to the right within a period. We will take a closer look at the exceptions later. This nucleus having a “tighter grip” on the outermost electrons also means that atomic radius decreases from the left to the right in a period.