Antimicrobial Properties of Chlorine and Alcohol Disinfectants
CHAPTER ONE
Objectives of the study
The main objective of the study is to assess the antimicrobial properties of chlorine and alcohol disinfectants.
This study compares the antimicrobial effect of a chlorine dioxide and a chlorine generating disinfectant on the contaminants commonly present on dental instruments and in the dental surgery.
Specifically, the study will;
- Determine the time duration of healing using animal models.
- Determine the efficacy of the chosen disinfectants.
- Identify the organism isolated from the wounds.
CHAPTER TWO
LITERATURE REVIEW
Concept of Chlorine
Chlorine is a chemical element with the symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the Pauling scale, behind only oxygen and fluorine.
The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, but not recognised as a fundamentally important substance. Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, romanized: khlôros, lit. ‘pale green’ based on its colour.
Because of its great reactivity, all chlorine in the Earth’s crust is in the form of ionic chloride compounds, which includes table salt. It is the second-most abundant halogen (after fluorine) and twenty-first most abundant chemical element in Earth’s crust. These crustal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater.
Elemental chlorine is commercially produced from brine by electrolysis. The high oxidising potential of elemental chlorine led to the development of commercial bleaches and disinfectants, and a reagent for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, and many intermediates for the production of plastics and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary. Elemental chlorine at high concentrations is extremely dangerous and poisonous for all living organisms, and was used in World War I as the first gaseous chemical warfare agent.
In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils as part of the immune response against bacteria.
Properties of Chlorine
Chlorine is the second halogen, being a nonmetal in group 17 of the periodic table. Its properties are thus similar to fluorine, bromine, and iodine, and are largely intermediate between those of the first two. Chlorine has the electron configuration [Ne]3s23p5, with the seven electrons in the third and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine. It is also a weaker oxidising agent than fluorine, but a stronger one than bromine. Conversely, the chloride ion is a weaker reducing agent than bromide, but a stronger one than fluoride. It is intermediate in atomic radius between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length. (Fluorine is anomalous due to its small size.)
All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: chlorine melts at −101.0 °C and boils at −34.0 °C. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure. The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow-green. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Specifically, the colour of a halogen, such as chlorine, results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital.[30] The colour fades at low temperatures, so that solid chlorine at −195 °C is almost colourless.
Like solid bromine and iodine, solid chlorine crystallises in the orthorhombic crystal system, in a layered lattice of Cl2 molecules. The Cl–Cl distance is 198 pm (close to the gaseous Cl–Cl distance of 199 pm) and the Cl···Cl distance between molecules is 332 pm within a layer and 382 pm between layers (compare the van der Waals radius of chlorine, 180 pm). This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable.
CHAPTER THREE
RESEARCH METHODOLOGY
Disinfectants
Chlorine dioxide (Aseptrol®) tablets manufactured by engelhard Aseptrol technology, USA and supplied by Waylor Trading and Logistics cc, South Africa were used in the study. A 48ppm solution of ClO2 was prepared by dissolving one 1.5g tablet of Aseptrol® in 2.5 litres of tap water according to the manufacturer’s instructions. The mixture was allowed to stand for 20 minutes. A further two- fold dilution of the disinfectant was prepared in tap water to give a final concentration of 24ppm solution.
Sodium dichloroisocyanurate (PreSept®, Johnson and Johnson, South Africa) solution was prepared by dissolving seven 2.5g tablets in one litre of tap water, according to the manufacturer’s instruction. This concentration is generally used for blood spillage and gave a final solution of 10 000ppm available chlorine. Both disinfectant solutions were prepared shortly before each experiment.
Test organisms and inocula
The antimicrobial properties of both disinfectants were tested against Staphylococcus aureus ATCC 29213, Pseudomonas aeruginosa ATCC 27853, Streptococcus mutans nCTC 1044, Candida albicans ATCC 90028, Bacillus subtilis (ATCC 6633) spores, Mycobacterium tuberculosis (ATCC 25177), Mycobacterium avium subsp. avium (ATCC 25291) and Hepatitis B virus (HBV). Stock cultures of S. aureus, P. aeruginosa, S. mutans, C. albicans and B. subtilis were stored in semisolid agar and subcultures were prepared as required. M. tuberculosis and M. avium subsp. avium were stored at -70oC in aliquots. Hepatitis B virus was obtained from the immunology laboratory, national Health Laboratory Services, Johannesburg, South Africa.
CHAPTER FOUR
RESULTS AND DISCUSSION
Seconds of exposure (Table 1) and B. subtilis spores in two to 2.5 minutes. The shelf-life test showed that chlorine dioxide killed all the test microorganisms within 30 seconds for up to 27 days whereas sodium dichloroisocyanurate was still effective after 37 days (Table 2).
CHAPTER FIVE
CONCLUSIONS
In conclusion this study has shown that a non-corrosive slow release chlorine dioxide and chlorine releasing sodium dichloroisocyanurate disinfectants are microbiocidal after 30 seconds exposure and sporicidal after two to three minutes in the presence of organic material. Both the disinfectant solutions were effective for 27 to 37 days if stored in screw cap bottles. They have the potential to be used in the dental setting as a surface disinfectant and a sterilant for semicritical heat sensitive instruments. Chlorine dioxide has an additional advantage because it is non-corrosive and the effective concentration is much lower than that required for sodium dichloroisocyanurate.
clinical significance
Chlorine dioxide has a broad-spectrum antimicrobial property and therefore it can be used in the dental settings for surfaces and heat sensitive instruments as a sterilant.
REFERENCES
- Awodele O, Emeka PM, Agbamuche HC,Akintowa A. The antimicrobial activities of somecommonly used disinfectants on Bacillus subtilis,Pseudomonas aeruginosa and Candida albicans.African J. of Biotechnology 2007; 6(8):987-990.
- Bloomfield S F. Chlorine and iodine formulations.In: Ascenzi J M, editor. Handbook of disinfectantsand antiseptics. New York, N. Y: Marcel Dekker,Inc., 1996; 33-158
- Chioma C. Okore, Ogechukwu N. Mbanefo,Bright C. Onyekwere, Simon C. Onyewenjo,Agaptus U. Ozurumba, Chinyere A.M. Abba-Father. Antimicrobial Efficacy of Selected Disinfectants, American Journal of Biology andLife Sciences, 2014, 2(2): 53-57.
- Coogan MM, Patel M, Mladenova D. efficacy of three surface disinfectants for dental radiographic films and gloves. J Dent 2006;32:385-9.
- Drake D, Villhauer AL. An in vitro comparative study determining bactericidal activity of stabilizing chlorine dioxide and other oral rinses. J Clin Dent 2011;22:1-5.
- Dychdala, G. R. (2001). Chlorine and chlorine compounds. In Disinfection, Sterilization and Preservation, 5th edn (Block, S. S., Ed.), pp. 135–58. Lippincott Williams & Wilkins, Philadelphia, PA, USA.
- Gottardi, W. & Karl, A. (1990). Chlorine covers on skin surfaces. I. Determination of the cover strength by the DPDcuvette-method. Zentralblatt fu¨r Hygiene und Umweltmedizin 190, 511–22.
- Gottardi, W. & Karl, A. (1991). Chlorine covers on skin surfaces. II. Parameters influencing the cover strength. Zentralblatt fu¨r Hygiene und Umweltmedizin 191, 478–93.
