Structure and Bonding Ionic bondsCovalent bonds Metallic bonds Chemical bonding involves either transferring or sharing electrons in the highest occupied energy level (shells) of atoms in order to achieve the electronic structure of a noble gas. P N N P P N N x x x Atoms can be represented as shown in this example; Mass Number 7 Li Atomic Number 3 Name of particle MassCharge Proton1+1 Neutron10 ElectronVery small Atoms of the same element can have different numbers of neutrons Atoms have an overall charge of 0 because the number of electrons is equal to the number of protons. When atoms form chemical bonds by transferring electrons, they form ions. Atoms that lose electrons become positively charged ions (+) because the total number of positively charged protons outnumber the number of negatively charged electrons. Atoms that gain electrons become negatively charged ions (-) because the total number of negatively charged electrons outnumber the number of positively charged protons. Ions have the electron structure of a noble gas in group 0 (take care as they do not ‘become’ a noble gas as the proton number remains unchanged). Elements in group 1 are called alkali metals and they all form ions with a charge of +1. They all react with non-metals to form ionic compounds. Elements in group 7 are called the halogens and they all form ions with a charge of -1. They all react with the alkali metals to form ionic compounds. An example is sodium chloride Sodium loses an electron to form Na+ Chlorine gains an electron to form Cl- The name of the ionic compound is Sodium chloride and the formula is NaCl - X X X X X X X X X X X X X X X X X X X X X X X X X X X + X When atoms share pairs of electrons, they form covalent bonds. These bonds between atoms are strong. Some covalently bonded substances consist of simple molecules such as H 2 (hydrogen), Cl 2 (chlorine), O 2 (oxygen), HCl (hydrochloric acid), H 2 O (water), NH 3 (ammonia) and CH 4 (methane). Others have giant covalent structures (macromolecules) such as diamond and silicon dioxide. Compounds with covalent bonds can be represented in 3 ways in order to show the shared pair of electrons. For example ammonia (NH 3 ) N H H H X X X XX O O O HH H N XOXO XOXO X O X HH H N Metals, such as copper and magnesium, consist of giant structures of atoms arranged in a regular pattern. The electrons in the highest occupied energy levels (outer shells) of metal atoms are delocalised so are free to move through the whole structure. This corresponds to a structure of positive ions with electrons between the ions holding them together by strong electrostatic forces

Structure and Properties Ionic structures Covalent structures Metallic bonds Substances that have simple molecular, giant ionic and giant covalent structures have very different properties. Ionic, covalent and metallic bonds are strong but the forces between molecules are weaker, e.g. in carbon dioxide and iodine. Substances that consist of simple molecules are gases, liquids and solids that have relatively low melting and boiling points. These substances have only weak forces between the molecules (intermolecular forces). It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils. Substances that consist of simple molecules do not conduct electricity because the molecules do not have an overall electric charge. Examples of simple molecules are oxygen, hydrogen, hydrochloric acid, ammonia, methane, carbon dioxide, water and chlorine. Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces in all directions between the oppositely charged ions. These compounds have high melting points and high boiling points because of the large amounts of energy needed to break the many strong bonds. When solid these structures do not conduct electricity because the ions cannot move. However when these substances melt or are dissolved in water they do conduct electricity because the ions are free to move. Atoms that share electrons can also form giant structures or macromolecules. Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures (lattices) of atoms. All the atoms in these structures are linked to other atoms by strong covalent bonds and so they have very high melting points. In diamond each carbon atom forms four covalent bonds with the other carbon atoms in a giant covalent structure so diamond is very hard. Diamond Graphite In graphite, each carbon atom bonds to three others forming layers. The layers are free to slide over each other because there are no covalent bonds between the layers and so graphite is soft and slippery. In graphite, one electron from each carbon atom is delocalised. These delocalised electrons allow graphite to conduct heat and electricity. Metals conduct heat and electricity because of the delocalised electrons in their outer shells. The layers of atoms in metals are able to slide over each other and so metals can be bent and shaped. Alloys are usually made from two or more different metals. The different sized atoms of the metals distort the layers in the structure, making it more difficult for them to slide over each other, and so make alloys harder than pure metals. Shape memory alloys can return to their original shape after being deformed, e.g. Nitinol used in dental braces Simple Covalent Molecules