Chemistry: Bonding

Ionic Bonding

Covalent Bonding

Desription

Ionic bonding form between metals and non-metals. It occurs when an atom with too many electrons in its outer shell passes electrons to an atom with too little electrons. The ultimate goal of this process is to get a full outer shell of electrons.In ionic bonding, it is the metal that loses the electrons to become a positively charged ion and the non-metal accepts the negatively charged electron to make it negatively charged. Ions have the same electronic structure as noble gases.It is the strong electrostatic forces of attraction that create the ions.

Example

This example shows the ionic bond of sodium chloride. The sodium atom loses an electron to the chloride atom to make sodium chloride where the sodium is positive and the chloride is negative.

Properties

  • High melting and boiling points - Ionic bonds are very strong - a lot of energy is needed to break them as there electrostatic forces are hard to weaken.This means that ionic compounds have high melting and boiling points.
  • Conductive when liquid - Ions are charged particles, but ionic compounds can only conduct electricity if their ions are free to move. Ionic compounds do not conduct electricity when they are solid - only when dissolved in water or melted.
  • Ionic solids are brittle and hard because the electrostatic attractions in the solid again hold the ions in definite positions. The electrostatic attractions must be overcome to move the ions. When the ions in the solid are shifted by some very strong force the positions of ions shift so that like charged ions are close together. This results in strong repulsions and the like charged ions move apart. The solid shatters and does not simply deform like a metal meaning that they are not malleable

Structure

The oppositely charged ions are arranged in a regular way to form giant ionic lattices. Ionic compounds often form crystals as a result.

Metallic

Desription

Exam ionic bond question

Explain why Metal Atoms form Positively Charged ions whereas non-metal atoms form negatively charged ions: When a metal reacts with a non-metal, each metal atom loses the electron or electrons in its outer shell leaving it with a full and stable electronic structure. The atom loses negative electrons but keeps and still has the same number of positive protons, leaving the metal ion with an overall positive charge. On the other hand, when the metal reacts with the non-metal, the non-metal atom gains electrons to give it a full outer shell. Consequently, this means that there are more negatively charged electrons than positively charged protons, showing why metal atoms form positively charged ions and non-metal atoms form negatively charged ions.

Metallic bonding is the force of attraction between delocalised electrons and the metal ions. It is the sharing of many detached electrons between many positive ions, where the electrons act as a "glue" giving the substance a definite structure. It is a strong electrostatic attraction. The outer electrons from each metal atom can easily move throughout the giant structure. The outer electrons form the sea of delocalised electrons surrounding the positively charged ions. The Strong electrostatic forces attraction between the negatively charged electrons and the positively charged ions bond the metal ions together. This can only happen between two or more metals.

Structure

The positive metal ions are surrounded by a sea of delocalised electrons. It is arranged in a giant lattice. The layers are weak as the electrostatic, intermolecular and Van Der Waals forces are weak.

Properties

  • High melting and boiling points
    Metallic bonds are strong and a lot of energy is needed to break them. This is why metals have high melting points and boiling points. The attraction between the metal ions and the delocalised electrons must be overcome to melt or to boil a metal. Some of the attractions must be overcome to melt a metal and all of them must be overcome to boil it. These attractive forces are strong, so metals have high melting and boiling points.
  • Electrically conductive
    Metals contain electrons that are free to move in the metal structure, carrying charge from place to place and allowing metals to conduct electricity well.The delocalised electrons are able to move through the metal structure. When a potential difference is applied, they will move together, allowing an electric current to flow through the metal.
  • Ductile and malleable
    Metals can be hammered and bent into different shapes, and drawn out into wires because of the layers being able to slide over each other.

Description

Covalent bonding is a type of atomic bonding that occurs between non-metals. It is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, Like all other bonds, the main aim is to maintain a stable structure.

Giant Covalent Structure

Simple Covalent structure

Properties

  • Low melting and boiling points - This is because the weak intermolecular forces break down easily.
  • Non-conductive - Substances with a simple molecular structure do not conduct electricity. This is because they do not have any free electrons or an overall electric charge.

Example

An example is carbon dioxide (CO2), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

Structure

Very strong van der waals forces but weak inter-molecular forces which hold them together.

Properties

  • Very high melting points - Substances with giant covalent structures have very high melting points, because a lot of strong covalent bonds must be broken. Graphite, for example, has a melting point of more than 3,600ºC.
  • Variable conductivity - Diamond does not conduct electricity. Graphite contains free electrons, so it does conduct electricity. Silicon is semi-conductive - that is, midway between non-conductive and conductive.

Structures

Giant covalent structures contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into giant regular lattices - extremely strong structures because of the many bonds involved.

Example

  • Graphite
    Graphite is a form of carbon in which the carbon atoms form layers. These layers can slide over each other, so graphite is much softer than diamond. It is used in pencils and as a lubricant. Each carbon atom in a layer is joined to only three other carbon atoms.
  • Diamond
    Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a giant covalent structure.
  • Silica, which is found in sand, has a similar structure to diamond. It is also hard and has a high melting point, but contains silicon and oxygen atoms, instead of carbon atoms.

Graphite

Diamond

Silica