1. Biochemistry Unit A Review on: Electronegativity, Bonding, Acids, Bases, Buffers, & Redox Reactions

3.  Hydrogen has an electronegativity of 2.1 Fluorine has an electronegativity of 4.0.  The difference between the two values will determine the type of bond that will form between the atoms.  4.0 – 2.1 = 1.9  If the difference is greater than 1.7 then the bond will be ionic.  If the difference is less than 1.7 then the bond will be covalent.  Therefore, the intramolecular bond that Hydrogen and Fluorine will form is ionic.

4.  Ionic bonds form when electrons are transferred between atoms.  By using the electronegativity values for each atom, you can determine the atom that will gain the electron and the atom that will lose the electron.  The atom with the larger electronegativity will take the electron and the atom with the lower electronegativity will lose its electron.  The atom that loses the electron will have a positive charge and forms a cation and the atom that gains the electron takes on a negative charge and forms an anion.

5.  Covalent bonds occur between two non-metals where the electronegativity difference is less than 1.7.  Covalent bonds involve the sharing of a pair of electrons between two atoms and create more stable compounds.  Covalent compounds are not soluble in water and will not conduct electricity. In addition, they generally have lower melting and boiling points and can exist in either a solid, liquid or gas state. Using the Lewis diagram for each atom will help you correctly draw the covalent compound that forms.  Covalent bonds can be single, double or triple.

6. Brønsted-Lowry Acids and Bases:  Acids can be defined as being proton donors  Bases are proton recipients  NH 3 is a base and HCl is an acid. If you consider the right side of this reaction reacting towards the left, you can consider NH 4 + an acid and Cl - a base.  So, every acid-base reaction can be considered as a reaction which produces a new acid and a new base:  These are called conjugate acid-base pairs. Note that the base paired to acid 1 is base 1. This are called a conjugate pair.  Also, the base paired with a strong acid is a weak base. The base paired with a weak acid is a strong base.

7. Lewis Acids and Bases:  A Lewis acid is defined as an electron-pair acceptor.  A Lewis base is an electron-pair donor.  For example, in the neutralization reaction  By eliminating the spectator ions, the reaction is essentially:  So the electron pair is donated by the OH-, thus it is the base, and the H+ accepts the electron pair and is the acid.

8. The pH scale: A Measure of [H + ]  For acids, the pH is below 7  For bases the pH is above 7  For neutral substances the pH is 7  Here are some sample values for pH. The pH of oven cleaners is about 14. Your blood maintains a pH between 7.35 and 7.45; if blood changes more than a few tenths of a pH unit from this normal range, the results could be fatal since too much acid or base interferes with the ability of the blood to pick up, carry and replace oxygen (denatures haemoglobin enzymes). Most plants grow best in soil with a pH value between 6 and 7; higher or lower values of pH prevent plants from absorbing nutrients from the soil. Most bacteria that cause food spoilage cannot grow in solutions having the low pH value of vinegar. Shampoos normally have a pH of about 8; your scalp has a pH of about 6.

9.  Many biological processes require specific pH levels. In order to maintain these vary narrow ranges, special types of solutions are needed: buffers. Buffers play an important role in biological systems maintaining the pH levels with an acceptable range to permit chemical reactions to proceed normally.  A buffer is a solution that resists changes in pH and maintains pH levels by taking up or releasing hydrogen ions or hydroxyl ions in a solution. It is made with a weak acid and a soluble salt containing the conjugate base of the weak acid or a weak base and a soluble salt containing the conjugate acid of the weak base.  A solution of the acid and its conjugate base form a compound that undergoes little change in pH when small quantities of strong acid or base are added. The weak acid reacts with any extra OH - ions put into the solution and the conjugate base reacts with any extra H + ions put into the solution.

10. Biological Example:  An example is carbonic acid and sodium bicarbonate system which occurs in the blood.  The pH of a healthy person's blood plasma is between 7.35 and 7.45. The blood's pH must be in the range of 7.0–7.8 in order to survive. If the pH is in the 7.0-7.3 range the person will feel tired, have trouble breathing, and may even be disoriented. If the pH of the blood is in the 7.5-7.8 range, the person will feel dizzy and rather agitated. If the pH is out of the 7.0– 7.8 range, oxygenated hemoglobin releases its oxygen which can eventually lead to death.

11.  Sliced fruit, such as apples, pears, and bananas, turn brown when exposed to air. The exposed flesh of the sliced fruit reacts with oxygen in the air to produce new brown coloured compounds. An enzyme in the fruit acts as a catalyst to speed up the browning reactions. The term oxidation can be used to describe this browning reaction of certain fruits when they react with oxygen. Try This…  Cut an apple in half. Take one half of the apple and rub lemon juice on the cut surface. Let both halves of the apple sit out exposed to the air, and periodically observe them for any changes in colour.  You may notice that the apple rubbed with the lemon juice did not turn brown as quickly as the untreated apple. The enzymes responsible for the browning reactions are sensitive to acids, so the acid in the lemon juice prevents the enzymes from working. Additionally, lemon juice contains vitamin C, which is highly reactive with oxygen. So the vitamin C in the lemon juice reacts with oxygen in the air before the fruit can.

12. REMEMBER: OIL RIG “Oxidation is loss” “Reduction is Gain” (of electrons)  As the study of chemistry advanced, the definitions of oxidation and reduction became much broader. According to modern chemical theory:  Oxidation is the loss of electrons.  Reduction is the gain of electrons.

13. Example 1:  If a piece of magnesium is placed in an aqueous solution of copper (II) sulfate, the magnesium displaces the copper in a single displacement reaction. This reaction is represented by the following equation:  Rewriting the equation to show the dissociated ions gives you:  The sulfate ion, SO 4 2-, is a spectator ion not involved in the chemical reaction. Cancelling out the sulfate ion gives you the net ionic equation for the reaction:  Looking at the net ionic equation, you can see the magnesium atom loses electrons to form a magnesium ion, while the copper ion gains electrons to form a copper atom.  In this reaction, the magnesium atom is oxidized because it lost electrons, while the copper (II) ion gained electrons and was reduced. Since oxidation and reduction both occur, the reaction is known as an oxidation- reduction, or redox reaction.

14. Oxidizing and Reducing Agents  Let's take another look at the net ionic equation of the redox reaction of magnesium and copper (II) sulfate: Take note of the following:  The electrons from the magnesium atom are transferred to the copper (II) ion.  The copper (II) ion is responsible for the oxidation of the magnesium atom. A reactant that oxidizes another reactant is called an oxidizing agent. An oxidizing agent accepts electrons.  The magnesium atom is responsible for the reduction of the copper (II) ion. A reactant that reduces another reactant is called a reducing agent. A reducing agent donates electrons in a redox reaction.  From the redox reaction of magnesium and copper (II) sulphate:

15. Oxidation and Reduction Half Reactions: To make it easier to follow the transfer of electrons in redox reactions, the oxidation and reduction reactions can be represented separately using half reactions.  A half reaction is a balanced chemical equation that shows the number of electrons being transferred. Since redox reactions involve both oxidation and reduction, two half reactions are required for redox reactions, one for oxidation and one for reduction.  From our earlier example, the reaction of magnesium with aqueous copper (II) sulphate can be represented by the net ionic equation:  Each neutral magnesium atom loses two electrons and is oxidized to form an Mg2+ ion.  The oxidation half reaction of this change is:  In the other half of the redox reaction, each copper (II) ion gains two electrons and is reduced to form a neutral copper atom.  The reduction half reaction of this change is:

16. Example 2:  The single displacement reaction of copper metal with silver nitrate produces solid silver and an aqueous solution of copper (II) nitrate.  Copper and silver nitrate react to produce a silver precipitate and a solution of copper (II) nitrate. The reaction of copper metal with silver nitrate can be represented by the balanced chemical equation below.  In this reaction, the nitrate ion, NO 3 1-, is a spectator ion, so removing it would leave the following net ionic equation:  In this reaction, each copper atom loses one electron and is oxidized to form a Cu1+ ion. This change is represented by the oxidation half reaction: Continued on next slide 

17. Example 2 continued…  Since the copper atom donates an electron, the copper atom is the reducing agent in this reaction.  Since our balanced chemical reaction involved two silver ions, each silver ion, Ag1+, gains an electron and is reduced to produce a silver atom. This change is represented by the reduction half reaction:  This half reaction could be simplified to show the reduction of one silver ion by dividing the reaction by two.  Since the silver ion accepts the electron, the silver ion is the oxidizing agent in this reaction. Don't forget to make sure all your redox reactions, including the half reactions which contain the electrons, are properly balanced. A proper balanced reaction will also balance charges: note how the half reactions above balance out the negatives and positives on each side of the reaction arrow.