Valency & bonding, Oxidation States and redox reaction

1. Valency & bonding, Oxidation States and redox reaction

2. Valency and Bonding • Most atoms consist of neutrons, protons (+) and electrons (-) • Neutrons and protons are contained within nucleus • Electrons (= number of protons or atomic number) are arranged in orderly shells outside. • Valence electrons are the outermost electrons of an atom • If electrons are lost, atom becomes +vely charged • If electrons are gained, atom becomes –vely charged

3. In the formation of ions, atoms of 2 elements undergo reduction and oxidation; 1 gains electrons and the other loses electrons.* • Metal element loses electron to gain a stable condition with no electrons in its outer shell. • The nonmetal steals electrons from the metal to complete its outer shell to 8 electrons, a stable configuration • E.g. chemical reaction between sodium and chlorine

4. Electron transfer during a chemical reaction, producing a sodium ion with an oxidation state of +1 and a chloride ion with an oxidation state of –1.

5. Oxidation state: A measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules: • (1) the oxidation state of a free element (uncombined element) is zero; • (2) for a simple (monoatomic) ion, the oxidation state is equal to the net charge on the ion; • (3) hydrogen has an oxidation state of 1 and oxygen has an oxidation state of -2 when they are present in most compounds.

6. (4) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion.

7. Element Oxidation number Group 1 elements in the Periodic Table +1 Group 2 elements in the Periodic Table +2 Al +3 Cl -1 O -2 Oxidation number are also called oxidation states. The oxidation number of an element in its uncombined state is zero. Example: Oxidation number of Cl2, O2 and Na is zero.

8. The sum of the oxidation numbers of all the elements in a compound is zero. For example, in FeCl3, the sum of the oxidation numbers of Fe and Cl = + 3 + 3 (-1) = 0. The sum of the oxidation numbers of all the atoms in a polyatomic ion equals the charge on the ion. For example, in SO42-, the sum of the oxidation numbers = + 6 + 4 (-2) = -2 (the charge on the sulphate ion).

9. Lewis structure Multiple oxidation states of chlorine due to sharing of electrons.

10. The valency/ oxidation number of an atom is determined by the number of electrons that it can take on, give up, or share with other atoms. • Valency is governed by the number of electrons in the outermost electronic shell of the atoms of that element (i.e. the valence electrons). • These valence electrons are given up to other atoms or are received from other atoms to make Ionic Bonds, or the valence electrons are shared with other atoms to make Covalent Bonds.

11. E.g. Water molecule • The Oxygen is two electrons short of a full outer shell. The Hydrogen atoms need to gain an extra electron to complete their outer shell which can hold two electrons. • Transferring electrons would require too much energy in this case. What happens is that the electrons are shared.

12. The Oxygen shares one of its electrons with the first Hydrogen atom and another with the second Hydrogen atom. • The Oxygen atom then has eight electrons orbiting it: equivalent to a full outer shell. • Each Hydrogen atom shares an electron with the Oxygen atom, thereby having two. • Molecules are discrete entities which have strong bonds between the atoms. This type of bond is called a Covalent Bond.

13. Chemical equations: weight relationship and conservation of mass and charge • Chemical equation • Mass must be conserved (total number of each kind of atom must be the same for both sides of equation) • Sum of charge on 1 side is equal to another • All formulas used must be correct H2 (g) + ½ O2 (g) → H2O (g) Reactants Products

14. Oxidation-reduction equations • An atom, molecule or ion is said to undergo oxidation when it loses an electron and to undergo reduction when it gains an electron • If electrons are equally shared by homonuclear atoms; no atom gains or loses electrons in the formation of molecule from its atom; the oxidation states = 0 (e.g. O2) • Heteronuclear polar covalent molecule has electrons which are shared unequally by atoms in the molecule (e.g. H2O) (see at http://web.jjay.cuny.edu/~acarpi/NSC/5-bonds.htm).

15. Oxidizing & Reducing Agent • Oxidizing agent*: Substances that cause some other species to be oxidized or to lose electrons. In order to cause the loss of electrons on some other species, the oxidizing agent must undergo reduction. • E.g. O(0), Cl(0), Fe(III), Cr(VI), Mn(IV), Mn(VII), N(V), N(III), S(0), S(IV), S(VI) • Reducing agent*: Substances that cause other species to be reduced or gain electrons. In order for reducing agents to cause the gaining of electrons on some other species they must undergo oxidation. • E.g. H(0), Fe(0), Mg(0), Fe(II), Cr(II), Mn(IV), N(III), Cl(-I), S(0), S(-II), S(IV)

16. Simple Oxidation –Reduction Reactions H20 + Cl20→ 2H+Cl- 4Fe0 + 3O20 → 2Fe23+O32- 2Fe2+ + Cl20 → 2Fe3+ + 2Cl- In each eq, the oxidizing agent gains the same number of e as are lost by the reducing agent

17. Complex Oxidation-Reduction Reactions - many oxidation-reduction reaction need acid or water • How to balance the eq of oxidation-reduction: • Write the skeleton eq. – molecular or ionic forms • Balance the eq with respect to the oxidation number change; balance the gain and lose of e • Complete the eq in the usual manner

18. Half Reaction • Half reactions are balanced oxidation-reduction reactions for a single element • They are not complete reactions because electrons are shown as of the reactants • A complete reaction is made by adding one half reaction to the reverse of another • Example: ClO3- (aq) + 2I- (aq) → Cl- (aq) + I2 (aq) ClO3- (aq) + 6H+ (aq) + 6e- → Cl- (aq) + 3H2O (l) 2I- (aq) → I2 (aq) + 6e- x 3: 6I- (aq) → 3I2 (aq) + 6e- Combine the two half-equations: CIO3- (aq) + 6H+ (aq) + 6I- (aq) → Cl-(aq) + 3I2 (aq) + 3H2O (l)