Topic 4 : Chemical Bonding & Structure

As a general rule, metals are on the left of the Periodic Table and non-metals are on the right-hand side

• Ionic bonds involve the transfer of electrons from a metallic element to a non-metallic element • Transferring electrons usually leaves the metal and the non-metal with a full outer shell • Metals lose electrons from their valence shell forming positively charged cations • Non-metal atoms gain electrons forming negatively charged anions • Cations and anions are oppositely charged and therefore attracted to each other • Electrostatic attractions are formed between the oppositely charged ions to form ionic compounds • This form of attraction is very strong and requires a lot of energy to overcome

Positive ion

Negative ion

• The ions form a lattice structure which is an evenly distributed crystalline structure • Ions in a lattice are arranged in a regular repeating pattern so that there are equal positive and negative charges • Therefore the final lattice is overall electrically neutral

• Sodium Chloride (salt) • Magnesium Oxide

• Strong • Brittle • High melting and boiling points • Soluble in water • Conduct electricity when molten or in solution

• Ionic compounds are strong • The strong electrostatic forces in ionic compounds keep the ions strongly together

• Ionic compounds are brittle as ionic crystals can split apart • When a layer of ions is pushed, it will repel with the ions on top causing the material to break

• Ionic compounds have high melting and boiling points • The strong electrostatic forces between the ions in the lattice act in all directions and keep them strongly together • Melting and boiling points increase with charge density of the ions due to the greater electrostatic attraction of charges • e.g. Mg2+O2- has a higher melting point than Na+Cl-

• Ionic compounds are soluble in water as they can form ion - dipole bonds • Basically, water has a positive and negative end so ionic ions attract and dissolve • Water has a positive and negative end because of the oxygen and hydrogen

• Ionic compounds only conduct electricity when molten or in solution • When molten or in solution, the ions can freely move around and conduct electricity • In the solid state they’re in a fixed position and unable to move around

• All metals form positive ions • There are some non-metal positive ions such as ammonium, NH4+, and hydrogen, H+

• The charge on the ions of the transition elements can vary which is why Roman numerals are often used to indicate their charge • Roman numerals are used in some compounds formed from transition elements to show the charge (or oxidation state) of metal ions • Eg. in copper (II) oxide, the copper ion has a charge of 2+ whereas in copper (I) nitrate, the copper has a charge of 1+

• The non-metals in group 15 to 17 have a negative charge and have the suffix ‘ide’ • Eg. nitride, chloride, bromide, iodide

• Also known as polyatomic ions • negative ions made up of more than one type of atom

• Ammonium : (NH4)+ • Hydroxide : (OH)- • Nitrate : (NO3)- • Sulfate : (SO4)2- • Carbonate : (CO3)2- • Hydrogen carbonate : (HCO3)- • Phosphate : (PO4)3-

• Covalent bonding occurs between two non-metals • A covalent bond involves the electrostatic attraction between nuclei of two atoms and the electrons of their outer shells • No electrons are transferred but only shared in this type of bonding • When a covalent bond is formed, two atomic orbitals overlap and a molecular orbital is formed • In a normal covalent bond, each atom provides one of the electrons in the bond. A covalent bond is represented by a short straight line between the two atoms, H-H • Covalent bonds should not be regarded as shared electron pairs in a fixed position; the electrons are in a state of constant motion and are best regarded as charge clouds

• Non-metals are able to share pairs of electrons to form different types of covalent bonds • Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas • This makes each atom more stable

The octet rule refers to the tendency of atoms to prefer to have eight electrons in the valence shell.

• In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell • Being able to accommodate more than 8 electrons in the outer shell is known as ‘expanding the octet rule’ • Accommodating less than 8 electrons in the outer shell means than the central atom is ‘electron deficient’

• The differences in Pauling electronegativity values can be used to predict whether a bond is covalent or ionic in character • A difference of less than around 1.0 in electronegativity values will be associated with covalent bonds, although between 1.0 and 2.0 can be considered polar covalent • A difference of more than 2.0 will be considered an ionic bond • All bonds have ionic or covalent character but some have less than others

• In diatomic molecules the electron density is shared equally between the two atoms • Eg. H2, O2 and Cl2 • Both atoms will have the same electronegativity value and have an equal attraction for the bonding pair of electrons leading to formation of a covalent bond

• In simple covalent bonds the two atoms involved share electrons • Some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom • An electron-deficient atom is an atom that has an unfilled outer orbital • In other words, both of the “shared” electron come from the same atom • This type of bonding is called dative covalent bonding or coordinate bond • An example is an ammonium ion