Tris buffer solution was prepared in the laboratory by dissolving 4.6g of boric acid, 6.5g of EDTA and 60.5g of trisma base in 1000cm³ of distilled water. Comparative study were done on the commercially produced tris buffer solution, using and laboratory produced tris buffer solution, using p^H meter for the p^Hand by carrying out genotype test for the efficacy of Tris buffer solution prepared in the laboratory was used as an electrolyte. This genotype test reduces the occurrence of sickle cell anemia. This expresses the efficacy rate and p^Hof laboratory produced Tris buffer solution as compared with that of commercially produced tris buffer solution. The result obtained for the p^H of laboratory produced tris buffer solution was found to be 9.0 while that of commercially produced tris buffer solution was found to be 8.9. the result obtained for the efficacy using genotype test showed that tris buffer solution produced in laboratory has more efficacy than that produced commercially, but the difference is not significant.

TABLE OF CONTENT

Certification i

Dedication ii

Acknowledgment iii

Abstract iv

Table of contents v

CHAPTER ONE

1.0 Introduction

1.1 Constituent of a buffer

1.2 Statement of problem

1.3 Objective of the research

1.4 Hypothesis

CHAPTER TWO

2.0 Literature review

2.1 Definition

2.2 Other types of buffer solutions

2.3 Importance and uses of buffers

2.4 Location of buffer in the body

2.5 Mechanism of action

2.6 What type of buffer is tris buffer solution

2.7 What is sickle cell anaemia

2.8 Physiological importance

2.9 Diagnostic use of buffer solution

2.10 Diagnostic significance

2.11 Storage/ shelf life

CHAPTER THREE

3.O Materials and methodology

3.1 Reagents

3.2 Apparatus/materials

3.3 Method of preparation

3.4 Comparative study

CHAPTER FOUR

4.0 RESULTS AND DISCUSSION

CHAPTER FIVE

5.0 CONCLUSIONS AND RECOMMENDATION

REFERNCES

CHAPTER ONE

1.0 INTRODUCTION

Martin et al (1989) defined buffer as a solution that has the ability to resist change in p^H, when small amounts of strong acids or bases are added to it. For example, when 0.01 mole of strong acids or base is added to distilled water, the p^H drops to 2 with the acids and rises to 12 with the base. If the same amount of acid or base is added to an acetic acid-sodium acetate buffer, the p^H may only change a fraction of a unit.

Buffers are important in many areas of chemistry. When the p^H must be controlled during the course of a reaction, the solutions are often buffered. This is often the case in biochemistry when enzymes or proteins are being studied. Our blood is often buffered to a p^H of 7.4. Variations of a few tenths of a p^H unit can cause illness or death. Acidosis is the condition when p^Hdrops too low. Alkalosis is the condition when p^H becomes higher than normal. Acidosis or alkalosis affects the functions of the heart.

In buffer solution two species are required one is capable of reaction with OH^- and the other will react with H₃O⁺. The two species must not react with each other. Many buffers are prepared by combing a weak acid and its conjugate (acetic acid and sodium acetate) or a weak base and its conjugate (ammonia and ammonium chloride). In general, the p^H range in which a buffer solution is effective is + or – one p^H unit on either side of the pka.

Buffering however is the tendency of a solution to resist change in p^H following addition of acid or base, according to Martin et al, 1989.

1.1 CONSTITUENT OF A BUFFER