Elm Chemistry Ch Pathology PDF

Reading Material for Medical Lab. Technician (Elementary chemistry and chemical pathology) Compiled By: Punjab Medical Faculty Specialized Healthcare & Medical Education Department Government of the Punjab...

Reading Material for Medical Lab. Technician (Elementary chemistry and chemical pathology) Compiled By: Punjab Medical Faculty Specialized Healthcare & Medical Education Department Government of the Punjab i  PREFACE This book is for a two years post matric teaching program of Medical Laboratory Technicians (MLTs) for the students of Allied Health Sciences. The purpose of this reading material is to provide basic education to the paramedics about elements, compounds including other chemistry topics. This book also contains knowledge of chemical pathology which includes the normal and pathological aspects of the biochemistry, interpretation. The reading material attempts to cover almost all the basic theoretical knowledge required by students about routine chemistry and chemical pathology so that they can give better performance in their work at Pathology laboratory. ii Elementary Chemistry and Chemical Pathology 1. Introduction to chemistry 1 2. Elements 5 3. Compounds and Mixtures 19 4. Units of measurement 29 5. Solutions 36 6. Acid, base and salt 46 7. Introduction of chemical pathology 68 8. Carbohydrates, lipids, proteins, enzymes and vitamins 86 9. Blood chemistry 111 10. Electrolytes and water 120 11.Significance of qualitative 129 iii  Chapter one Introduction to chemistry Objective After studying this chapter students will be able to demonstrate scientific understandings of the structure of matter and of its physical and chemical transformations. Students will be able to apply appropriate theories to predict chemical structure, reactivity, and physical properties. iv  CHAPTER 1 INTRODUCTION TO CHEMISTRY Chemistry deals with reality’s most basic elements, from particles to atoms to molecules, chemistry is also known as the central science. Chemistry is also concerned with the utilization of natural substances and the creation of artificial ones. Cooking, fermentation, glass making, and metallurgy are all chemical processes that date from the beginnings of civilization. Today, vinyl, Teflon, liquid crystals, semiconductors, and superconductors represent the fruits of chemical technology. The 20th century saw dramatic advances in the comprehension of the marvelous and complex chemistry of living organisms, and a molecular interpretation of health and disease holds great promise. Modern chemistry, aided by increasingly sophisticated instruments, studies materials as small as single atoms and as large and complex as DNA (deoxyribonucleic acid), which contains millions of atoms. New substances can even be designed to bear desired characteristics and then synthesized. The rate at which chemical knowledge continues to accumulate is remarkable. Over time more than 8,000,000 different chemical substances, both natural and artificial, have been characterized and produced. Chemistry, deals with the properties, composition, and structure of substances (defined as elements and compounds), the transformations they undergo, and the energy that is released or absorbed during these processes. Every substance, whether naturally occurring or artificially produced, consists of one or more species of atoms that have been identified as elements. Although these atoms, in turn, are composed of elementary particles, they are the basic building blocks of chemical substances; there is no quantity of oxygen, mercury, or gold, for example, smaller than an atom of that substance. Chemistry, therefore, is concerned not with the subatomic domain but with the properties of atoms and the laws governing their combinations and how the knowledge of these properties can be used to achieve specific purposes. Sitting between biology and physics, the field of chemistry is sometimes called the central science. This branch of science deals with the most basic elements of reality, such as fundamental particles, or the complex world of living organisms, but the in-between world of atoms, molecules and chemical processes. Chemistry is the study of matter, analyzing its structure, properties and behavior to see what happens when they change in chemical reactions. As such, it can be considered a branch of physical science, alongside astronomy, physics and earth sciences including geology. An important area of chemistry is the understanding of atoms and what determines how they react. It turns out, reactivity is often largely mediated by the electrons that orbit atoms and the way these are exchanged and shared to create chemical bonds. 1  1.1, Branches of Chemistry Chemistry has now split into many branches. For instance, analytical chemists might measure the traces of compounds in ancient pottery to discern what people were eating thousands of years ago. 1.1.1, Biochemistry is the study of the chemical processes that take place in living organisms, for instance in farming, and on the effect the resulting produce will have on our body’s metabolism. It examines the processes that occur in molecules and cells, the communications between cells and the relation of a cell's structure to its functions. Biochemistry has broad applications in medicine, particularly in studying the causes and cures of viruses, in nutrition and agriculture. Research in biochemistry may cover anything from basic cellular processes up to understanding disease to develop better treatments. 1.1.2, Organic chemistry, the study of compounds which contain carbon, connects up molecules in new ways to build and analyse an array of materials, from drugs to plastics to flexible electronics. 1.1.3, Inorganic chemistry: is the study of chemicals that are not primarily based on carbon. Inorganic chemicals are commonly found in rocks and minerals. One area of inorganic chemistry deals with the design and properties of materials involved in energy and information technology. Inorganic chemistry has many practical applications, including 2 fertilizers production, surfactants and pigments. Inorganic chemistry is applied in the field of chemical and pharmaceutical production. 1.1.4, Physical chemistry involves looking at chemistry through the lens of physics to study changes in pressure, temperatures and rates of conversion, for example, as substances react. 1.1.5, Analytical chemistry: a branch of chemistry that uses chemical analysis to define types of matter and determine their quantities. There are qualitative and quantitative methods of chemical analysis. Qualitative methods look at the presence of different matters, whereas quantitative methods determine how much of a chemical exists within matter. There are many applications of analytical chemistry, including food and drug safety, environmental regulations, medical diagnosis, and forensic science. Most of the materials that occur on Earth, such as wood, coal, minerals, or air, are mixtures of many different and distinct chemical substances. Each pure chemical substance (e.g., oxygen, iron, or water) has a characteristic set of properties that gives it its chemical identity. Iron, for example, is a common silver-white metal that melts at 1,535° C, is very malleable, and readily combines with oxygen to form the common substances hematite and magnetite. The detection of iron in a mixture of metals, or in a compound such as magnetite, is a branch of analytical chemistry called qualitative analysis. Measurement of the actual amount of a certain substance in a compound or mixture is termed quantitative analysis. Quantitative analytic measurement determines that iron makes up 72.3 percent, by mass, of magnetite. Some very simple qualitative tests reveal the presence of specific chemical elements in very smaller amounts. The yellow color imparted to a flame by sodium. Such analytic tests have allowed chemists to identify the types and amounts of impurities in various substances and to determine the properties of very pure materials. 3 Sample Question; Define chemistry and name the branches of Chemistry? References 1; Thomas Holme, J. Chem. Educ. 2022, 99, 10, 3353–3354 Publication Date:October 11, 2022 2 Luxford, C. J.; Holme, T. A. What do Conceptual Holes in Assessment Say about the Topics We Teach in General Chemistry. J. Chem. Educ. 2015, 92 (6), 993– 1002, DOI: 10.1021/ed500889j 3 Kovarik, M. L.; Galarreta, B. C.; Mahon, P. J.; McCurry, D. A.; Gerdon, A. E.; Collier, S. M.; Squires, M. E. Survey of the Undergraduate Analytical Chemistry Curriculum. J. Chem. Educ. 2022, 99 (6), 2317– 2326, DOI: 10.1021/acs.jchemed.2c0009 4 Hunter, K. H.; Rodrigues, J-M. G.; Becker, N. M. A Review of Research on the Teaching and Learning of Chemical Bonding. J. Chem. Educ. 2022, 99 (7), 2451– 2464, DOI: 10.1021/acs.jchemed.2c00034 5 Dood, A. J.; Watts, F. M. Mechanistic Reasoning in Organic Chemistry: A Scoping Review of How Students Describe and Explain Mechanisms in the Chemistry Education Research Literature. J. Chem. Educ. 2022, 99 (8), 2864– 2876, DOI: 10.1021/acs.jchemed.2c00313 4  Chapter 2 Elements Objective of the chapter, After studying his chapter, the students will learn how the chemical elements are arranged on the Periodic Table. Students will be able to arrange the elements on the Periodic Table. Students will learn about some of the important chemical and physical properties of the elements. 5  Chemical element, Any substance that cannot be decomposed into simpler substances by ordinary chemical processes. Elements are the fundamental materials of which all matter is composed. At present there are 118 known chemical elements. About 20 percent of them do not exist in nature (or are present only in trace amounts) and are known only because they have been synthetically prepared in the laboratory. Of the known elements, 11 are gases under ordinary conditions, 2 are liquids (two more, cesium and gallium, melt at about or just above room temperature), and the rest are solids. Elements can combine with one another to form a wide variety of more complex substances called compounds. The number of possible compounds is almost infinite; perhaps a million are known, and more are being discovered every day. When two or more elements combine to form a compound, they lose their separate identities, and the product has characteristics quite different from those of the constituent elements. The gaseous elements hydrogen and oxygen, for example, with quite different properties, can combine to form the compound water, which has altogether different properties from either oxygen or hydrogen. Water clearly is not an element because it consists of, and actually can be decomposed chemically into, the two substances hydrogen and oxygen; these two substances, however, are elements because they cannot be decomposed into simpler substances by any known chemical process. Most samples of naturally occurring matter are physical mixtures of compounds (seawater), the most common of which is sodium chloride, or table salt. Atoms of elemental substances are themselves complex structures composed of more fundamental particles called protons, neutrons, and electrons. Within an atom, a small nucleus, which contains both protons and neutrons, is surrounded by cloud, of electrons. The fundamental properties of these subatomic particles are their weight and electrical charge. Whereas protons carry a positive charge and electrons a negative one, neutrons are electrically neutral. The diameter of an atom (about 10−8 centimetre) is 10,000 times larger than that of its nucleus. Neutrons and protons, which are collectively called nucleons, have relative weights of approximately one atomic mass unit, whereas an electron is only about 1/2000 as heavy. Because neutrons and protons occur in the nucleus, virtually all of the mass of the atom is concentrated there. The number of protons in the nucleus is equivalent to the atomic number of the element. The total number of protons and neutrons is called the mass number. Because the atom itself is electrically neutral, the atomic number represents not only the number of protons, or positive charges, in the nucleus but also the number of electrons, or negative charges, in the extra nuclear region of the atom. The chemical characteristics of elements are intimately related to the number and arrangement of electrons in their atoms. Thus, elements are completely distinguishable from each other by their atomic numbers. The realization that such is the case leads to another definition of an element, namely, a substance, all atoms of which have the same atomic number. Not all the atoms present have the same atomic weight, even though they all have the same atomic number. Such a situation can occur only if the atoms have different numbers of neutrons in their nuclei. Such groups of atoms—with the same atomic number but with different relative weights—are called isotopes. The number of isotopic forms that a naturally occurring element possesses ranges from one (e.g., fluorine) to as many as ten (e.g., tin); 6 most of the elements have at least two isotopes. The atomic weight of an element is usually determined from large numbers of atoms containing the natural distribution of isotopes, and, therefore, it represents the average isotopic weight of the atoms constituting the sample. 2.1, Atom, The basic building block of all matter and chemistry. Atoms can combine with other atoms to form molecules but cannot be divided into smaller parts by ordinary chemical processes. Most of the atom is empty space. The rest consists of three basic types of subatomic particles: protons, neutrons, and electrons. The protons and neutrons form the atom’s central nucleus. The ordinary hydrogen atom is an exception; it contains one proton but no neutrons. As their names suggest, protons have a positive electrical charge, while neutrons are electrically neutral—they carry no charge, the nucleus has a positive charge. Circling the nucleus is a cloud of electrons, which are negatively charged. Like opposite ends of a magnet that attract one another, the negative electrons are attracted to a positive force, which binds them to the nucleus. The nucleus is small and dense compared with the electrons, which are the lightest charged particles in nature. The electrons circle the nucleus in orbital paths called shells, each of which holds only a certain number of electrons. An ordinary, neutral atom has an equal number of protons in the nucleus and electrons surrounding the nucleus. Thus the positive and negative charges are balanced. Some atoms, however, lose or gain electrons in chemical reactions or in collisions with other particles. Ordinary atoms that either gain or lose electrons are called ions. If a neutral atom loses an electron, it becomes a positive ion. If it gains an electron, it becomes a negative ion. These basic subatomic particles—protons, neutrons, and electrons—are themselves made up of smaller substances, such as quarks and leptons. More than 90 types of atoms exist in nature, and each kind of atom forms a different chemical element. They are ranked in order of their atomic number (the total number of protons in its nucleus) in a chart called the periodic table. Accordingly, because an atom of iron has 26 protons in its nucleus, its atomic number is 26 and its ranking on the periodic table of chemical elements is 26. Because an ordinary atom has the same number of electrons as protons, an element’s atomic number also tells how many electrons its atoms have, and it is the number and arrangement of the electrons in their orbiting shells that determines how one atom interacts with another. The key shell is the outermost one, called the valence shell. If this outermost shell is complete, or filled with the maximum number of electrons for that shell, the atom is stable, with little or no tendency to interact with other atoms. But atoms with incomplete outer shells seek to fill or to empty such shells by gaining or losing electrons or by sharing electrons with other atoms. This is the basis of an atom’s chemical activity. Atoms that have the same number of electrons in the outer shell have similar chemical properties. Electrons contribute only a tiny part to the mass of the atomic structure, however, they play an important role in the chemical reactions that create molecules. For most purposes, the atomic weight can be thought of as the number of protons plus the number of neutrons. Because the number of neutrons in an atom can vary, there can be several different atomic 7 weights for most elements. Atoms having the same number of protons but different numbers of neutrons represent the same element and are known as isotopes of that element. An isotope for an element is specified by the sum of the number of protons and neutrons. For example, the following are two isotopes of the carbon atom: Carbon 12 is the most common, non-radioactive isotope of carbon. Carbon 14 is a less common, radioactive carbon isotope. 8 Figure: Structure of atom 9  2.2, Periodic Table of the Elements The Periodic table of elements, simply called a Periodic Table, is a systematic arrangement of 118 known chemical elements. These chemical elements are organized in order of increasing atomic number, or the number of protons in an atom’s nucleus, which typically corresponds with increasing atomic mass, from left to right and top to bottom. The horizontal rows from left to right are called periods while the vertical columns from top to bottom are called groups in a periodic table. The periodic table is an essential aspect of Chemistry because it is an arrangement of all the known elements and therefore provides information about elements and their relation with one another to use reference. For example, properties of particular elements such as their mass, electron number, electron configuration, and their unique chemical properties. The 118 elements are arranged in 7 periods and 18 groups as shown below. Further, the elements are divided into different blocks. 2.2.1 s-block elements: The first and the second group elements that have the last electron filled in the s-subshell are called s-block elements. The elements included in the s-block are Alkali metals and Alkaline Earth Metals. 2.2.2 p-block elements: The elements included in groups 13 to 17 are p-block elements. These elements have the last electron filled in their p-subshell. The elements included in this block from different groups are termed as Boron Family (Group 13), Carbon Family (Group 14), Nitrogen Family (Group 15), Oxygen Family (Group 16), and Fluorine Family (Group 17). 2.2.3 d-block elements: The elements present in groups 3 to 12 are d-block elements. These elements have the last electron filled in their d-subshell. d-block elements are also called transition elements as they have partially filled d-orbitals in their ground state. 2.2.4 f-block elements: These elements have the last electron filled in their f- subshell. Such elements are present in Lanthanides and Actinide groups. Here is the table representing 118 elements of the periodic table. The elements listed are arranged according to the increasing order of atomic number and their respective atomic weight, symbols, density, and electronegativity. 10  2.3, Metal, any of a class of substances characterized by high electrical and thermal conductivity and high reflectivity of light. Metals are solids at room temperature (except for mercury), and are usually malleable (can be rolled into sheets) and ductile (can be drawn into wires). Metals are usually separated into the main group metals in Groups IA - VA and the transition metals in Groups IB - VIIIB. Approximately three-quarters of all known chemical elements are metals. The most abundant varieties in the Earth’s crustare aluminum, iron, calcium, sodium, potassium, and magnesium. The vast majority of metals are found in ores (mineral-bearing substances), but a few such as copper, gold, platinum, and silver frequently occur in the free state because they do not readily react with other elements. Metals are usually crystalline solids. In most cases, they have a relatively simple crystal structure distinguished by a close packing of atoms and a high degree of symmetry. Typically, the atoms of metals contain less than half the full complement of electrons in their outermost shell. Because of this characteristic, metals tend not to form compounds with each other. They do, however, combine more readily with nonmetals (e.g., oxygen and sulfur), which generally have more than half the maximum number of valence electrons. Metals differ widely in their chemical reactivity. The most reactive include lithium, potassium, and radium, whereas those of low reactivity are gold, silver, palladium, and platinum. The high electrical and thermal conductivities of the simple metals (i.e., the non-transition metals of the periodic table) are best explained by reference to the free-electron theory. According to this concept, the individual atoms in such metals have lost their valence electrons to the entire solid, and these free electrons that give rise to conductivity move as a group throughout the solid. In the case of the more complex metals (i.e., the transition elements), conductivities are better explained by the band theory, which takes into account not only the presence of free electrons but also their interaction with so- called d electrons. 11 The mechanical properties of metals, such as hardness, ability to resist repeated stressing, ductility, and malleability, are often attributed to defects or imperfections in their crystal structure. The absence of a layer of atoms in its densely packed structure, for example, enables a metal to deform plastically, and prevents it from being brittle. 2.4, The Non-Metals are a group of elements located on the right side of the periodic table (except for hydrogen, which is on the top left). These elements are distinctive in that they typically have low melting and boiling points, don't conduct heat or electricity very well, and tend to have high ionization energies and electronegativity values. They also don't have the shiny "metallic" appearance associated with the metals. While the metals are malleable and ductile, the nonmetals tend to form brittle solids. The nonmetals tend to gain electrons readily to fill their valence electrons shells, so their atoms often form negative-charged ions. There are 7 elements that belong to the nonmetals group: Hydrogen (sometimes considered an alkali metal), Carbon, Nitrogen, Oxygen, Phosphorus, Sulfur, Selenium Even though there are only 7 elements within the nonmetals group, two of these elements (hydrogen and helium) make up about 98% of the mass of the universe.1Nonmetals form more compounds than metals. Living organisms consist mainly of nonmetals. 2.5, MetalloidS, is a term used to describe a chemical element that forms a simple substance having properties intermediate between those of a typical metal and a typical nonmetal. The term is normally applied to a group of between six and nine elements (boron, silicon, germanium, arsenic, antimony, tellurium, bismuth, polonium, astatinefound near the center of the P-block or main block of the periodic table. There is no single property which can be used to unambiguously identify an element as a metalloid. Since most metalloids tend to display semiconducting properties, the class might reasonably be extended to also include gray silicon (which, unlike white silicon, is a semiconductor rather than a metal) and the graphite form of carbon (which, unlike the diamond form, is a semimetal rather than an insulator). Chemically, metalloids correspond to atoms having intermediate electronegativities and an ability to display a range of both positive and negative oxidation states in their compounds. It can be noted that all seven of these elements can be found on the regular periodic table in a diagonal region of the p-block which extends from boron (which is placed on the upper left) to astatine (which is placed on the lower right). Some periodic tables have a dividing line between metals and nonmetals, and below this line, the metalloids can be found. 2.6, Named Families 2.6.1, Alkali Metals Group IA, belong to the s-block elements occupying the left most side of the periodic table. Alkali metals readily lose electrons, making them count among the most reactive elements on earth. In general, ‘alkali’ refers to the basic or alkaline nature of their metal hydroxides. The compounds are called alkali metals because when they react with water, they usually form alkalies which are nothing but strong bases that can easily neutralise acids. They occupy the first column of the periodic table. Alkali elements 12 are Lithium(Li), Sodium(Na), Potassium (K), Rubidium (Ru), Cesium (Cs) and Francium (Fr), occupying successive periods from first to seven. Francium is a radioactive element with a very low half-life. However, the main reason why hydrogen (H) is not considered an alkali metal is that it is mostly found as a gas when the temperature and pressure are normal. Hydrogen can show properties or transform into an alkali metal when it is exposed to extremely high pressure. 2.6.2, Alkaline earth metals are the elements that correspond to group 2 of the modern periodic table. This group of elements includes beryllium, magnesium, calcium, strontium, barium, and radium. The elements of this group are quite similar in their physical and chemical properties. For example, all alkaline earth metals are silvery-white-coloured solids under standard conditions. They are also highly shiny and are quite reactive. Since the alkaline earth metals have a completely full s-orbital in their respective valence shells, they tend to readily lose two electrons to form cations with a charge of +2. Thus, the most common oxidation state exhibited by alkaline earth metals is +2. The alkaline earths are the elements located in Group IIA of the periodic table. This is the second column of the table. 2.6.3, Halogens, any of the six nonmetallic elements that constitute Group 17 Group VIIA of the periodic table. The halogen elements are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Ts). They were given the name halogen, from the Greek roots hal- (“salt”) and -gen (“to produce”), because they all produce sodium salts of similar properties, of which sodium chloridetable salt, or halite is best known. Because of their great reactivity, the free halogen elements are not found in nature. In combined form, fluorine is the most abundant of the halogens in Earth’s crust. The percentages of the halogens in the igneous rocks of Earth’s crust are 0.06 fluorine, 0.031 chlorine, 0.00016 bromine, and 0.00003 iodine. Astatine and tennessine do not occur in nature, because they consist of only short-lived radioactive isotopes. The halogen elements show great resemblances to one another in their general chemical behaviour and in the properties of their compounds with other elements. There is, however, a progressive change in properties from fluorine through chlorine, bromine, and iodine to astatine. Fluorine is the most reactive of the halogens. Chlorine is the best known of the halogen elements. The free element is widely used as a water-purification agent, and it is employed in a number of chemical processes. Table salt, sodium chloride, of course, is one of the most familiar chemical compounds. Fluorides are known chiefly for their addition to public water supplies to prevent tooth decay, but organic fluorides are also used as refrigerants and lubricants. Iodine is most familiar as an antiseptic, and bromine is used chiefly to prepare bromine compounds that are used in flame retardants and as general pesticides. In the past ethylene dibromide was extensively used as an additive in leaded gasoline. 13  2.6.4, Group VIIIA of the periodic table are the noble gases or inert gases: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). The name comes from the fact that these elements are virtually unreactive towards other elements or compounds. They are found in trace amounts in the atmosphere (in fact, 1% of the atmosphere is argon); helium is also found in natural gas deposits. In their elemental form at room temperature, the Group 8A elements are all colorless, odorless, monatomic gases. They traditionally have been labeled Group 0 in the periodic table because for decades after their discovery it was believed that they could not bond to other atoms; that is, that their atoms could not combine with those of other elements to form chemical compounds. Their electronic structures and the finding that some of them do indeed form compounds has led to the more appropriate designation, Group 18 (VIIIA). Under standard conditions for temperature and pressure, all the noble gases exist in the gaseous phase. They are known to possess extremely low chemical reactivity (hence the name inert gas). This is because all the noble gases have stable electronic configurations. This is the reason why noble gases do not form molecules easily and are mostly found as mono-atomic gases. 14  2.7, Valence, The property of an element that determines the number of other atoms with which an atom of the element can combine. The term is used to express both the power of combination of an element in general and the numerical value of the power of combination. Characteristic valences for the elements were measured in terms of the number of atoms of hydrogen with which an atom of the element can combine or that it can replace in a compound. It became evident, however, that the valences of many elements vary in different compounds. Valence electrons are the s and p electrons in the outermost shell. The electrons present in the inner shell are core electrons. When we study and observe the atom of an element, we come across tiny subatomic particles called valence electrons. Valence electrons are all arranged in different orbitals or shells and are mostly negatively charged particles. Further, these electrons are responsible for interaction between atoms and the formation of chemical bonds. However, not all electrons are associated with the atom. Only the electrons present in the outermost shell can participate in the formation of a chemical bond or a molecule. Such type of electrons is called valence electrons. 2.8, Symbols Each element has a name. Each element name is abbreviated as usually the first letter of the element’s name, while the second letter is some other letter from the name. Some elements have symbols that derive from earlier, mostly Latin names, so the symbols may not contain any letters from the English name. Table lists the names and symbols of some of the most familiar elements. Element Names and Symbols Aluminum Al magnesium Mg Argon Ar manganese Mn Arsenic As mercury Hg* Barium Ba Neon Ne Bismuth Bi Nickel Ni Boron B nitrogen N Bromine Br oxygen O Calcium Ca phosphorus P Carbon C platinum Pt Chlorine Cl potassium K* Chromium Cr Silicon Si Copper Cu* Silver Ag* Fluorine F Sodium Na* Gold Au* strontium Sr Helium He Sulfur S Hydrogen H Tin Sn* Iron Fe tungsten W† 15  Iodine I uranium U Lead Pb* Zinc Zn Lithium Li zirconium Zr *The symbol comes from the Latin name of element. †The symbol for tungsten comes from its German name— wolfram. 2.9, DESCRIPTION OF IMPORTANT ELEMENTS most living matter consists primarily of the bulk elements: oxygen, carbon, hydrogen, nitrogen, and sulfur, the building blocks of the compounds. these five elements also constitute the bulk of our diet. six other elements, sodium, magnesium, potassium, calcium, chlorine, and phosphorus are often referred to as macrominerals because they provide essential ions in body fluids and form the major structural components of the body. in addition, phosphorus is a key constituent of both dna and rna: the genetic building blocks of living organisms. some of these important elements are listed below: 2.9.1, CARBON, our body is made up of 18% carbon. sugar, proteins, vitamins, etc all are made up of carbon. apart from this carbon is used in medicinal industries for making drugs and medicines. carbon dioxide gas we release during respiration is used by plants for the process of photosynthesis. 2.9.2, HYDROGEN, it is used in the synthesis of water which is essential for life. it is used in making fertilizers, ammonia, etc. other than this hydrogen is used as rocket fuel to create strong explosions. 2.9.3, OXYGEN it is essential for respiration, circulation, and many more physiological processes. we cannot live without oxygen. 2.9.3, CHLORINE, It is used as a disinfectant for water. industrial and sewage waste help in sanitization. it is used as a bleaching agent. 2.9.4, SULPHUR, it is used in the synthesis of sulfuric acid and in making insecticides and fungicides. 2.9.5, CALCIUM, makes strong bones and teeth. it is important for physiological processes like the movement of muscles, for nerves, carry signals from the brain to body parts. calcium carbonate is used in the cement industry in the synthesis of cement. 16  2.9.6, IRON is an important element of blood and helps in the transportation of food and minerals apart from this iron is used in industries in anufacturing structural elements for building, machinery and tools, and many more. 2.9.7, PHOSPHORUS, it is used in making matches, fertilizers, phosphorus bronze, detergents, and many more. 2.9.8, NITROGEN, is an essential element in the synthesis of proteins, in medicine manufacturing, food packing, and preservation, etc, 2.9.9, SODIUM, it helps in synthesis of many reagents in chemical industries (sodium hydroxide) naoh, sodium borohydrides (nabh4) etc. apart from this sodium salts are used in medicinal industries. the salt we consume in our food is nacl (sodium chloride). other than these elements many other elements like magnesium, zinc, neon, and helium are also used in our daily life. 17 Sample Question; describe the definition of valence, metalloids, symbols and periodic table? Refernces 1 Philos Trans A Math Phys Eng Sci. 2015 Mar 13; 373(2037): 20140182. 2 Fraga CS. 2005. Relevance, essentiality and toxicity of trace elements in human health. Mol. Aspects Med. 26, 235–244. 3 International Year of the Periodic Table of Chemical Elements (IYPT) 2019: Planning, Coordination, and Implementation, IUPAC Project 2018-005-2-020. Co-chaired by Jan 4 Reedijk and Natalia Tarasova (https://iupac.org/project/2018-005-2-020).. Stephen A. Matlin, Goverdhan Mehta, Henning Hopf, Alain Krief, First published: 28 March 2019 https://doi.org/10.1002/ejic.201801409 C 5 E.I. Hamilton Volume 3, Issue 1, September 1974, Pages 3-85. ∗ 18  Chapter 3 Compounds and Mixtures Learning Objective of the chapter: After studying this chapter students will be able to know what makes compounds different from elements. Students will be able to identify compounds by their names and properties. 19  3.1, Chemical compound, any substance composed of identical molecule consisting of atoms of two or more chemical elements. The fundamental principle of the science of chemistry is that the atoms of different elements can combine with one another to form chemical compounds. Methane, for example, which is formed from the elements carbon and hydrogen in the ratio four hydrogen atoms for each carbon atom, is known to contain distinct CH4 molecules. The formula of a compound such as CH4 indicates the types of atoms present. Water, which is a chemical compound of hydrogen and oxygen in the ratio of two hydrogen atoms for every oxygen atom, contains H2O molecules. Sodium chloride is a chemical compound formed from sodium (Na) and chlorine (Cl) in a 1:1 ratio. Although the formula for sodium chloride is NaCl. The substances mentioned above exemplify the two basic types of chemical compounds: molecular (covalent) and ionic. Methane and water are composed of molecules; that is, they are molecular compounds. Sodium chloride, on the other hand, contains ions; it is an ionic compound. In fact, there are millions of chemical compounds known, and many more millions are possible. Most substances found in nature—such as wood, soil, and rocks—are mixtures of chemical compounds. These substances can be separated into their constituent compounds by physical methods, which are methods that do not change the way in which atoms are aggregated within the compounds. Compounds can be broken down into their constituent elements by chemical changes. A chemical change (that is, a chemical reaction) is one in which the organization of the atoms is altered. An example of a chemical reaction is the burning of methane in the presence of molecular oxygen (O 2) to form carbon dioxide (CO2) and water. CH4 + 2O2 → CO2 + 2H2O Chemical compounds at ordinary temperatures and pressures, some are solids, some are liquids, and some are gases. The colours of the various compounds span those of the rainbow. Some compounds are highly toxic to humans, whereas others are essential for life. Substitution of only a single atom within a compound may be responsible for changing the colour, odour, or toxicity of a substance. Out of this great diversity, classification systems have been developed. An example is the classification of the compounds as molecular or ionic. Compounds are also classified as organic or inorganic. Because of the great variety of ways that carbon can bond with itself and other elements, there are more than nine million organic compounds. 20  3.2, Mixtures A mixture is a compound that is made up of two more chemical compounds or substances that do not combine together chemically. It is the physical combination of two or more substances that are able to retain their individual identities while they are mixed to form solutions, suspensions, or colloids. Physical means can be used to separate them. A solution of salt and water, a combination of sugar and water, various gases, air, and so on are examples. The different components of any combination do not unite through any chemical changes. As a result, the components retain their distinct characteristics. In addition, unlike in a compound, the components in a mixture do not combine chemically to produce new material. Instead, they just mix and maintain their original characteristics. Properties of Mixtures, All the components or substances in a mixture retain their original physical properties. The mixture can be separated into its components physically by using some techniques Components in a mixture may or may not be in a fixed proportion and can vary in quantity. Examples of Mixtures 3.2.1 Smog is a mixture of Smoke and Fog. 3.2.2 Cement is a mixture of Sand, Water and Gravel. 3.2.3 Sea Water is a mixture of Water and Salt. 3.2.4 Soil is a mixture of Minerals, Air, Organic materials, Water, and Living Organisms. 3.2.5 Blood is a mixture of Plasma, White Blood Cells, Red Blood Cells, and Platelets. 3.2.6, Gasoline is a mixture of Hydrocarbons, Petroleum, and Fuel Additives. 3.3, Evaporation, a process by which an element or compound transitions from its liquid state to its gaseous state below the temperature at which it boils; in particular, the process by which liquid water enters the atmosphere as water vapour in the water cycle. 21 In order for a liquid molecule to escape into the gas state, the molecule must have enough kinetic energy to overcome the intermolecular attractive forces in the liquid. A given liquid sample will have molecules with a wide range of kinetic energies. Liquid molecules that have this certain threshold kinetic energy escape the surface and become vapor. As a result, the liquid molecules that remain now have lower kinetic energy. As evaporation occurs, the temperature of the remaining liquid decreases. On a hot day, the water molecules in perspiration absorb body heat and evaporate from the surface of your skin. The evaporation process leaves the remaining perspiration cooler, which in turn absorbs more heat from your body. A given liquid will evaporate more quickly when it is heated. This is because the heating process results in a greater fraction of the liquid's molecules having the necessary kinetic energy to escape the surface of the liquid. 3.4, Distillation, involves the conversion of a liquid into vapour that is subsequently condensed back to liquid form. When steam from a kettle becomes deposited as drops of distilled water on a cold surface. Distillation is used to separate liquids from nonvolatile solids, as in the separation of gasoline, kerosene, and lubricating oil from crude oil. Other industrial applications include the processing of chemical products as formaldehyde and phenol and the desalination of seawater. 22 Distillation refers to the selective boiling and subsequent condensation of a component in a liquid mixture. It is a separation technique that can be used to either increase the concentration of a particular component in the mixture or to obtain pure components from the mixture. The process of distillation exploits the difference in the boiling points of the components in the liquid mixture by forcing one of them into a gaseous state. It is important to note that distillation is not a chemical reaction but it can be considered as a physical separation process. Role of Raoult’s Law and Dalton’s Law The temperature at which the vapor pressure of a liquid becomes equal to the pressure of the surrounding area is known as the boiling point of that liquid. At this temperature point, the liquid is converted into its vapor form via the formation of vapor bubbles at its bulk. For a mixture of liquids, the distillation process is dependent on Dalton’s law and Raoult’s law. As per Raoult’s law, the partial pressure of a single liquid component in an ideal liquid mixture equals the product of the vapor pressure of the pure component and its mole fraction. According to Dalton’s law of partial pressures, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of all the constituent gases. When a mixture of liquids is heated, the vapor pressure of the individual components increases, which in turn increases the total vapor pressure. Therefore, the mixture cannot have multiple boiling points at a given composition and pressure. Types of Distillation Simple distillation 23  Fractional distillation Steam distillation Vacuum distillation Air-sensitive vacuum distillation Short path distillation Zone distillation 3.5, Filtration, Filtration is technically defined as the process of separating suspended solid matter from a liquid, by causing the latter to pass through the pores of a membrane, called a filter. The process in which solid particles in a liquid or gaseous fluid are removed by the use of a filter medium that permits the fluid to pass through but retains the solid particles. Either the clarified fluid or the solid particles removed from the fluid may be the desired product. Other media, such as electricity, light, and sound, also can be filtered. The basic requirements for filtration are: (1) a filter medium; (2) a fluid with suspended solids; (3) a driving force such as a pressure difference to cause fluid to flow; and (4) a mechanical device (the filter) that holds the filter medium, contains the fluid, and permits the application of force. Filtration also plays a role in water treatment. The process of filtration can become a costly process when it comes to water treatment and water purification. But filters have enough advantages to be used as a mechanism of water treatment or purification. Filtration Methods 3.5.1, General Filtration: The most basic form of filtration is using gravity to filter a mixture. The mixture is poured from above onto a filter medium (e.g., filter paper) and gravity pulls the liquid down. The solid is left on the filter, while the liquid flows below it. 3.5.2, Vacuum Filtration: A Büchner flask and hose are used to create a vacuum to suck the fluid through the filter (usually with the aid of gravity). This greatly speeds the separation and can be used to dry the solid. A related technique uses a pump to form a pressure difference on both sides of the filter. Pump filters do not need to be vertical because gravity is not the source of the pressure difference on the sides of the filter. 3.5.3, Cold Filtration: Cold filtration is used to quickly cool a solution, prompting the formation of small crystals. This is a method used when the solid is initially dissolved. A common method is to place the container with the solution in an ice bath prior to filtration. 3.5.4, Hot Filtration: In hot filtration, the solution, filter, and funnel are heated to minimize crystal formation during filtration. Stemless funnels are useful because there is less surface area for crystal growth. This method is used when crystals would clog the funnel or prevent crystallization of the second component in a mixture. 24  3.6, Sedimentation Process Sedimentation is the process of allowing particles in suspension in water to settle out of the suspension under the effect of gravity. The particles that settle out from the suspension become sediment, and in water treatment is known as sludge. When a thick layer of sediment continues to settle, this is known as consolidation. When consolidation of sediment, or sludge, is assisted by mechanical means then this is known as thickening.Sedimentation is the process of separating small particles and sediments in water. This process happens naturally when water is still because gravity will pull the heavier sediments down to form a sludge layer. However, this action can be artificially stimulated in the water treatment process. This mechanical assistance is called thickening. 25  Uses of Sedimentation Sedimentation also helps to determine the medical conditions of a person. Sedimentation rate of RBC’s is an example. The sedimentation rate is performed by measuring how long it takes red blood cells (RBCs) to settle in a test tube. As time passes, RBC’s start to separate from the other plasma contents, and they settle down at the bottom and serum will be formed above. The sedimentation rate or the ESR (erythrocyte sedimentation rate) is measured simply by recording how far the top of the Red Blood Cell layer has fallen (in millimetres) from the top of the serum layer in one hour. Water treatment plants use the method of sedimentation to filter out unwanted particles from unclean water. Filtering through several layers of sand and soil, allowing certain sizes of particles to pass through. A separatory funnel is used to separate immiscible liquids. When two immiscible liquids are placed in a separator funnel, two layers are seen. The denser solvent will be the bottom layer. Most halogenated solvents are denser than water, most non halogenated solvents are less dense than water. If you are not sure which layer is which, add a drop of water and see which layer it joins. 26  3.7, Magnetic separation is a process used to separate materials based on their magnetic properties. Essentially, magnetic materials can be separated from non-magnetic materials by using a magnetic field. The application of a magnetic field in the separation process is initiated by using a magnet or an electromagnet, which produces a magnetic field in the desired separation region. Within this field, materials possessing magnetic properties are drawn towards the magnet’s source. This attraction is often due to the material’s magnetic orientations of its atoms synchronize with the external magnetic field. Conversely, non- magnetic materials remain unaffected by the magnet, allowing for their easy separation as they don’t adhere to the magnet. Following the separation, magnetic and non-magnetic materials are gathered. This separation process can be conducted in both dry and wet environments. There are various devices designed for this purpose, such as: 3.7.1, Magnetic Drum Separator: Used for the automatic separation of magnetic particles from raw materials. 3.7.2, Magnetic Roll Separator: Uses strong rare-earth magnets for the purification of products. 3.7.3, Magnetic Pulley: Often installed at the end of conveyor belts to extract metal contaminants. 27  Sample Question; Describe the types of compounds and the procedure of filtration? References; 1 : J. Chem. Educ. 2012, 89, 7, 832–833, Publication Date:May 3, 2012, What Are Elements and Compounds? 2 Cite this: J. Chem. Educ. 2007, 84, 5, 880, Publication Date:May 1, 2007 A2: Element or Compound? 3 Miguel Reina, Hervé This, Antonio Reina. Improving the Understanding of Chemistry by Using the Right Words: Why Is Talking about Compounds so Messy?. Journal of Chemical Education 2024, 101 (1) , 39-48. https://doi.org/10.1021/acs.jchemed.3c00557 4 Juan Quílez. A categorisation of the terminological sources of student difficulties when learning chemistry. Studies in Science Education 2019, 55 (2) , 121- 167. https://doi.org/10.1080/03057267.2019.1694792 28  Chapter 4 Units of Measurements Objectives. This chapter relates various units to appropriate quantities for measurement. It shows that compound units can be produced by multiplying and by dividing the units associated with quantities and understand that the product of a unit and a value is constant for a measurement. 29  Measurements and Units Measurements provide the macroscopic information that is the basis of most of the hypotheses, theories, and laws that describe the behavior of matter and energy in both the macroscopic and microscopic domains of chemistry. Every measurement provides three kinds of information: a number (quantitative observation), a unit (describes how it was measured), and the degree of reliability (uncertainty of the measurement). While the number and unit are explicitly represented when a quantity is written, the uncertainty is an aspect of the measurement result. The number in the measurement can be represented in different ways, including decimal form and scientific notation. For example, the maximum takeoff weight of a Boeing 777-200ER airliner is 298,000 kilograms, which can also be written as 2.98 × 105 kg. The mass of the average mosquito is about 0.0000025 kilograms, which can be written as 2.5 × 10−6 kg. 4.1, Units, such as liters, pounds, and centimeters, are standards of comparison for measurements. When we buy a 2-liter bottle of a soft drink, we expect that the volume of the drink was measured, so it is two times larger than the volume that everyone agrees to be 1 liter. Without units, a number can be meaningless, confusing, or possibly life threatening. Suppose a doctor prescribes phenobarbital to control a patient’s seizures and states a dosage of “100” without specifying units. Not only will this be confusing to the medical professional giving the dose, but the consequences can be dire: 100 mg given three times per day can be effective as an anticonvulsant, but a single dose of 100 g is more than 10 times the lethal amount. We usually report the results of scientific measurements in SI units, an updated version of the metric system, using the units listed in Table 1. Other units can be derived from these base units. The standards for these units are fixed by international agreement, and they are called the International System of Units or SI Units (from the French, Le Système International d’Unités). Table: Base Units of the SI System Property Measured Name of Unit Symbol of Unit Length Meter M mass Kilogram Kg time Second S temperature Kelvin K electric current Ampere A amount of substance Mole Mol luminous intensity Candela Cd 30 Sometimes we use units that are fractions or multiples of a base unit. these fractions or multiples are always powers of 10. Fractional or multiple SI units are named using a prefix and the name of the base unit. For example, a length of 1000 meters is also called a kilometer because the prefix kilo means “one thousand,” which in scientific notation is 103 (1 kilometer = 1000 m = 103 m). SI also provides a series of prefixes that can be attached to the units, creating units that are larger or smaller by powers of 10. Common prefixes and their multiplicative factors are listed in Table. The base unit kilogram is a combination of a prefix, kilo- meaning 1,000 ×, and a unit of mass, the gram.) Some prefixes create a multiple of the original unit: 1 kilogram equals 1,000 grams, and 1 megameter equals 1,000,000 meters. Other prefixes create a fraction of the original unit. Thus, 1- centimeter equals 1/100 of a meter, 1millimeter (ml) equals 1/1,000 of a meter, 1 microgram equals 1/1,000,000 of a gram, and so forth. Table: Prefixes Used with SI Units Prefix Abbreviation Multiplicative Multiplicative Factor in Factor Scientific Notation giga- G 1,000,000,000 × 109 × mega- M 1,000,000 × 106 × kilo- K 1,000 × 103 × deca- D 10 × 101 × deci- D 1/10 × 10−1 × centi- C 1/100 × 10−2 × milli- M 1/1,000 × 10−3 × micro- µ* 1/1,000,000 × 10−6 × nano- N 1/1,000,000,000 × 10−9 × *The letter µ is the Greek lowercase letter for m and is called “mu,” which is pronounced “myoo.” 31  4.2, SI Base Units There are several SI units used in physics that are used to express the different quantities. The quantities can be classified into two groups i.e. base units and derived units. These are the fundamental units and are considered as the building blocks of the system. All the other units are derived from the SI Base units. One of the examples is that the SI unit of mass is kilogram. There are 7 SI base units. The seven units along with their SI unit and symbol are given below: 4.2.1, Unit of length, meter (m): Meter is the SI unit of length and is defined by taking the fixed value of the speed of light in vacuum. It is expressed as m.s -1. It is defined as the distance light in a vacuum travel in 1/299,792,458 of a second. A meter is about 3 inches longer than a yard. one meter is about 39.37 inches or 1.094 yards. Longer distances are often reported in kilometres (1 km = 1000 m = 10 3 m), whereas shorter distances can be reported in centimetres (1 cm = 0.01 m = 10−2 m) or millimetres (1 mm = 0.001 m= 10−3 m). 4.2.2, Unit of mass, kilogram (kg): Kilogram is the SI unit of mass and is defined by taking the fixed value of the Planck constant. It is expressed as kg.m 2.s-1. A kilogram was originally defined as the mass of a liter of water (a cube of water with an edge length of exactly 0.1 meter). It is now defined by a certain cylinder of platinum-iridium alloy. Any object with the same mass as this cylinder is said to have a mass of 1 kilogram. One kilogram is about 2.2 pounds. The gram (g) is exactly equal to 1/1000 of the mass of the kilogram (10−3 kg). 4.2.3, Unit of time, second (s): Second is the SI unit of time and is defined by taking the fixed value of Cesium frequency. It is expressed as s1. Small- and large- time intervals can be expressed with the appropriate prefixes; for example, 3 microseconds = 0.000003 s = 3 × 10−6 and 5 Mega seconds = 5,000,000 s = 5 × 106 s. Alternatively, hours, days, and years can be used. 4.2.4, Unit of electric current, ampere (A): Ampere is the SI unit of electric current and is defined by taking the fixed value of the elementary charge. 4.2.5, Unit of thermodynamic temperature, Kelvin (K): Kelvin is the SI unit of thermodynamic temperature and is defined by taking the fixed value of Boltzmann constant k = 1.380649×10-23. The degree Celsius (°C) is also allowed in the SI system, with both the word “degree” and the degree symbol used for Celsius measurements. Celsius degrees are the same magnitude as those of kelvin, but the two scales place their zeros in different places. Water freezes at 273.15 K (0 °C) and boils at 373.15 K (100 °C). 4.2.6, Unit of luminous intensity, candela (cd): Candela is the SI unit of luminous intensity and is defined by the fixed value of the luminous efficacy. 32  4.2.7, Unit of the amount of substance; mole (mol): Mole is the SI unit of the amount of substance and is defined by the fixed value of Avogadro constant NA. One mole contains 6.02214076×1023 elementary entities and is expressed as mol- 1. 4.3, Derived SI Units, SI derived units are units of measurement derived from the seven SI base units specified by the International System of Units (SI). They can be expressed as a product (or ratio) of one or more of the base units, possibly scaled by an appropriate power of exponentiation. The derived units are unlimited as they are formed by different operations on the base units. For derived units, the dimensions are expressed in terms of the dimensions of the base units. The derived units might also be expressed with the combination of base and derived units. 4.4, Volume Volume is the measure of the amount of space occupied by an object. The standard SI unit of volume is defined by the base unit of length. The standard volume is a cubic meter (m3), a cube with an edge length of exactly one meter. A more commonly used unit of volume is derived from the decimetre (0.1 m, or 10 cm). A cube with edge lengths of exactly one decimeter contains a volume of one cubic decimeter (dm3). A liter (L) is the more common name for the cubic decimeter. One liter is about 1.06 quarts. A cubic centimeter (cm3) is the volume of a cube with an edge length of exactly one centimeter. The abbreviation cc (for cubic centimeter) is often used by health professionals. A cubic centimeter is also called a milliliter (mL) and is 1/1000 of a liter. 4.5, The Celsius Scale The Celsius scale of the metric system is named after Swedish astronomer Anders Celsius (1701-1744). The Celsius scale sets the freezing point and boiling point of water at 0oC and 100oC, respectively. The distance between those two points is divided into 100 equal intervals, each of which is one degree. Another term sometimes used for the Celsius scale is "centigrade" because there are 100 degrees between the freezing and boiling points of water on this scale. However, the preferred term is "Celsius". 33  4.6, The Kelvin Scale It is based on molecular motion, with the temperature of 0K0K, also known as absolute zero, being the point where all molecular motion ceases. The freezing point of water on the Kelvin scale is 273.15K273.15K, while the boiling point is 373.15K373.15K. Notice that there is no "degree" used in the temperature designation. Converting between Scales The Kelvin is the same size as the Celsius degree, so measurements are easily converted from one to the other. The freezing point of water is 0°C = 273.15 K; the boiling point of water is 100°C = 373.15 K. The Kelvin and Celsius scales are related as follows: K=°C+273.15 °C=5÷9×(°F−32) °F=9÷5×(°C) +32 There is only one temperature for which the numerical value is the same on both the Fahrenheit and Celsius scales: −40°C = −40°F.s Each measurement has an amount, a unit for comparison, and an uncertainty. Measurements can be represented in either decimal or scientific notation Scientists primarily use the SI (International System) or metric systems. We use base SI units such as meters, seconds, and kilograms, as well as derived units, such as liters (for volume) and g/cm3 (for density). In many cases, we find it convenient to use unit prefixes that yield fractional and multiple units, such as microseconds (10 −6 seconds) and megahertz (106 hertz), respectively. 34 Sample Question; describe the units of measurements, Mass , Temperature? What are SI units of Measurement? References; 1; J Am Med Inform Assoc. 1999 Mar-Apr; 6(2): 151–161. Units of Measure in Clinical Information Systems. Gunther Schadow, MD, Clement J. McDonald, MD, Jeffrey G. Suico, MD, Ulrich Föhring, MD, and Thomas Tolxdorff, PhD. 2. ISO 2955: Information Processing: Representation of SI and Other Units in Systems with Limited Character Sets. Geneva, Switzerland: International Organization for Standardization, 1983. 3, Warren W. Tryon, Vol. 17, No. 3 (Summer 1996), pp. 213-227 (15 pages), Measurement Units and Theory Construction. 4; Luca Mari, Charles Ehrlich and Leslie Pendrill, Measurement units as quantities of objects or values of quantities: A discussion, August 2018 Metrologia 55(5). 35  Chapter 5 Solutions Objectives This chapter will help students distinguish between heterogeneous and homogeneous mixtures, different solute and solvent combinations, compare the properties of suspensions and solutions. It helps in comparing the effects of temperature and pressure on solubility. the mass of solute and volume of solvent, and the concentration of a solution, and to determine the amount of solute in a given amount of solution. 36  5.1, Solutions Solutions have two components; one is solvent and the other is solute. Aqueous solutions are those where the solvent is water. Sugar in water, carbon dioxide in water, etc. are examples. Non-aqueous solutions do not use water as a solvent. The component that dissolves the other component is called the solvent. The component dissolved in the solvent is called solute. Generally solvent is present in a major proportion compared to the solute. The amount of solute is lesser than the solvent. The solute and solvent can be in any state of matter i.e. solid, liquid, or gas. Solutions that are in the liquid state consist of a solid, liquid, or gas dissolved in a liquid solvent. Alloys and air are examples of solid and gaseous solutions, respectively. Examples The following examples illustrate solvent and solute in some solutions. 5.1.1 Air is a homogeneous mixture of gases. Here both the solvent and the solute are gases. 5.1.2 Sugar syrup is a solution where sugar is dissolved in water using heat. Here, water is the solvent and sugar is the solute. 5.1.3 Tincture of iodine, a mixture of iodine in alcohol. Iodine is the solute whereas alcohol is the solvent. 5.2, Types of Solution Liquid solutions, such as sugar in water, are the most common kind, but some solutions are gases or solids. 37 Any state of matter (solid, liquid, or gas) can act both as a solute and as a solvent. 5.2.1, Aqueous Solutions; Aqueous Solutions contain water as the solvent. Different solutes can be dissolved in water to form such solutions, such as salt water, sugar water or carbon dioxide in water. 5.2.2 Non-Aqueous Solutions; Non-Aqueous Solutions do not contain water as the solvent. The solvent could be other liquids such as ether, petrol, carbon tetrachloride, etc. Some examples of non-aqueous solutions are sulphur in carbon disulphide, naphthalene in benzene, etc. 5.2.3 Saturated Solutions; A solvent can dissolve some particular types of solutes in it. The maximum amount of solute that can be dissolved in a solvent at a specified temperature can be termed a saturated solution. A solution cannot dissolve any more solute further upon reaching saturation. The undissolved substances remain at the bottom. The point at which the solute stops dissolving in the solvent is termed the saturation point. 5.2.4 Unsaturated Solutions; The amount of solute that is contained in lesser amounts than the maximum value, that is before the solution reaches the saturation level is called an unsaturated solution. No remaining substances leave at the bottom, that is, all the solute is dissolved in the solvent. An unsaturated solution is basically a chemical solution which has a solute concentration lesser than its corresponding equilibrium solubility. 5.2.5 Supersaturated Solutions; The amount of solute contained in the solution exceeds the maximum amount of solute. The solution has already reached and crossed the saturation point. The solute is dissolved into the solution forcefully by raising the temperature or pressure of the solution. The solute particles on further dissolve, crystal out in the bottom of the container by the method called crystallization. 5.2.6 Concentrated Solutions; A concentrated solution contains large quantities of solute in the given solvent to form a solution. Some examples of concentrated solutions are mango juice, brine solution or dark colour tea. 5.2.7 Dilute Solutions; A dilute solution contains small quantities of solute in the given large quantity of solvent to form a solution. Some examples of dilute solutions are salt solutions or light color tea. 5.2.8 Isotonic Solution; The solution contained in the beaker has a higher concentration of solute in it. As a result of this, the water emerges from the cell and into the solution contained in the beaker. 5.2.9 Hypertonic Solution; Hypertonic solutions contain the same concentration of solute in them. The water moves across the cell from the solution in the beaker in both directions. 5.2.10 Hypotonic Solution; There is a lower concentration of solute in the solution contained in the beaker. As a result, water goes into the cell which causes the cells to swell up and eventually burst. 5.3, Properties of Solutions Properties of the solutions are as follows: 5.3.1 A solution is a homogeneous mixture. 5.3.2 The constituent particles of a solution are smaller than 10-9 metres in diameter. 38  5.3.3 Constituent particles of a solution cannot be seen by naked eyes. 5.3.4 Solutions do not scatter a beam of light passing through it. So, the path of the light beam is not visible in solutions. 5.3.5. Solute particles cannot be separated by filtration. 5.3.6 Solute or solvent particles do not settle down when left undisturbed. 5.3.7 Solutions are stable at a given temperature. 5.4, Solubility The maximum amount of solute that can be dissolved in a known quantity of solvent at a certain temperature is called solubility. The factors affecting solubility vary on the state of the solute: Liquids in Liquids Solids in Liquids Gases In Liquids Factors Affecting Solubility: The solubility of a substance depends on the physical and chemical properties of that substance. In addition to this, there are a few conditions which can manipulate it. Temperature, pressure and the type of bond and forces between the particles are few among them. 5.4.1 Temperature: By changing the temperature, we can increase the soluble property of a solute. Generally, water dissolves solutes at 20° C or 100° C. Sparingly soluble solid or liquid substances can be dissolved completely by increasing the temperature. But in the case of gaseous substance, temperature inversely influences solubility i.e. as the temperature increases gases expand and escapes from their solvent. 5.4.2 Forces and Bonds: The type of intermolecular forces and bonds vary among each molecule. The chances of solubility between two unlike substances are more challengeable than the like substances. For example, water is a polar solvent where a polar solute like ethanol is easily soluble. 39  5.4.3 Pressure: Gaseous substances are much influenced than solids and liquids by pressure. When the partial pressure of gas increases, the chance of its solubility is also increased. A soda bottle is an example of where CO 2 is bottled under high pressure. 5.4.4 Solubility of Liquids in Liquids Solubility is the new bond formation between the solute molecules and solvent molecules. In terms of quantity, solubility is the maximum concentration of solute that dissolves in a known concentration of solvent at a given temperature. Solutes are categorized into highly soluble, sparingly soluble or insoluble. If a concentration of 0.1 g or more of a solute can be dissolved in a 100ml solvent, it is said to be soluble. While a concentration below 0.1 g is dissolved in the solvent it is said to be sparingly soluble. Thus, it is said that solubility is a quantitative expression and is expressed by the unit gram/liter (g/l). 5.4.5 Solubility of Solids in Liquids Let us understand the process by which a solid dissolves in a solvent. Once a solid solute is added to a solvent, the solute particles dissolve in the solvent and this process is known as dissolution. Solute particles in the solution collide with each other and some of these particles get separated out of the solution, this process is called crystallization. A state of dynamic equilibrium is established between these two processes and at this point, the number of solute molecules entering the solution becomes equal to the number of particles leaving the solution. As a result, the concentration of the solute in the solution will remain constant at a given temperature and pressure. 5.4.6 Solubility of Gases In Liquids The gas solubility in liquids is greatly affected by temperature and pressure as well as the nature of the solute and the solvent. There are many gases that readily dissolve in water, while there are gases that do not dissolve in water under normal conditions. Oxygen is only sparingly soluble in water while HCl or ammonia readily dissolves in water. Henry’s Law 40 gives a quantitative relation between pressure and gas solubility in a liquid. It states that: The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas present above the surface of liquid or solution. P = KHx Where, p = partial pressure of the gas x = mole fraction of the gas solution KH = Henry’s law constant 5.5 Concentration of a Solution The amount of solute in a given solution is called the concentration of a solution. The proportion of solute and solvent in solutions is not even. Depending upon the proportion of solute, a solution can be: Diluted Concentrated Saturated The concentration of solution = Amount of solute Amount of solution 5.5.1 Molarity (Molar Concentration): Molarity (M) is defined as a number of moles of solute dissolved in one litre (or one cubic decimetre) of the solution. The unit of molarity is mol L-1 0r mol dm-3 or M. Number of moles of a substance can be found using the formula Molarity changes with temperature because volume changes with temperature. Molarity can be expressed as Decimolar = M/10 (0.1 M) Semimolar = M/2 (0.5 M) Pentimolar = M/5 (0.2 M) Centimolar = M/100 (0.01 M) milimolar = M/1000 (0.001 M). 41  5.5.2 Molality: Molality (m) is defined as a number of moles of solute expressed in kg dissolved in one kg of solvent, Molality has no unit. Molality is a better way of expressing concentration than molarity because there is no term of volume of solvent is involved. The volume of the solvent depends on the temperature of the solvent. Thus, there is no effect of the change of temperature on the molality. Molality is related to solubility as 5.5.3 Normality: Normality (N) is defined as gram-equivalent of solute dissolved in one litre (or one cubic decimetre) of the solution, Unit of molarity is N. 42  5.5.4 Formality: Formality is the number of formula mass in gram present per litre of a solution. If the formula mass of solute is equal to its molar mass, then the formality is equal to molarity. The formality of a solution depends on temperature. This concept is used in the case of ionic substances. A mole of an ionic compound is called formole and its molarity is called formality. Thus, the formality of a solution may be defined as a number of moles of ionic solute present in one litre of the solution. 43 Sample Question; What are the types of a solution? Describe factors affecting solubility? References; 1; Ketan T. Savjani, Anuradha K. Gajjar, * and Jignasa K. Savjani, ISRN Pharm. 2012; 2012: 195727. Published online 2012 Jul 5. doi: 10.5402/2012/195727. Drug Solubility: Importance and Enhancement Techniques. 2 Vemula VR, Lagishetty V, Lingala S. Solubility enhancement techniques. International Journal of Pharmaceutical Sciences Review and Research. 2010;5(1):41–51. 3; Narinder Singh, Amar paul Singh, January 2021, International Journal of Pharmaceutical Chemistry and Analysis 7(4):166-171. Solubility: An overview 44  Chapter 6 Acids, Bases and Salts Objectives This chapter introduces students to tell the difference between an acid and a base. Students will learn to define and identify both types of substances and explain how they differ. They will also discover why this type of information is important to know. Acids and bases are important to many chemical processes: maintaining a stable internal environment in the human body, baking a delicious cake, or determining whether a lake can support aquatic life. Reactions involving acids and bases can be described through the transfer of protons – single H+ ions. 45  6.1 Acids Acids are characterized by their ability to donate hydrogen ions (H+) when dissolved in water. An acid is a substance whose water solution tastes sour, turns blue litmus red and neutralizes bases. According to Liebig, acids are compounds which contain hydrogen that can be replaced by metals. Acids have pH of less than 7. Acids react with bases to form salts and water. Acids can be found naturally in many foods and beverages, including citrus fruits, vinegar, and fermented products, and they are also used in various industrial processes. Acids can be classified into two categories: 6.1.1 organic acids, which are derived from living organisms, 6.1.2 inorganic acids, which are derived from non-living sources. 6.2 Chemical Properties of Acid Acid has various chemical properties few of the following chemical properties of acids include, 6.3 Reaction of acids with metals: When an acid reacts with a metal, it produces hydrogen gas and the corresponding salt. Metal + Acid → Salt + Hydrogen When hydrochloride acid combines with zinc metal, it produces hydrogen gas and zinc chloride. Zn + 2HCl → ZnCl2 + H2 When acids react with metal carbonates, they produce carbon dioxide gas and salts as well as water. Metal carbonate + Acid → Salt + Carbon dioxide + Water When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water. Na2CO3 + 2HCl → 2NaCl + H2O + CO2 Reaction of acid with hydrogen carbonates (bicarbonates): When acids react with metal hydrogen carbonates, they produce carbon dioxide gas, salt, and water. Acid + Metal hydrogen carbonate → Salt + Carbon dioxide + Water Sulfuric acid gives sodium sulfate, Carbon dioxide gas and water when it reacts with sodium bicarbonate. 2NaHCO3 + H2SO4 → NaCl + CO2 + H2O 46  6.4 Types of Acids On the basis of thei Occurrence acid are subdivided into two categories Natural Acids Mineral Acids Natural Acids; Natural acids also known as organic acids, are acids derived from natural sources. For example, Methanoic acid (HCOOH), Acetic acid (CH3COOH), Oxalic acid (C2H2O4), etc. Mineral Acids These are created from minerals. Inorganic acids, man-made acids, and synthetic acids are all examples of Mineral Acids. For example, Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid (H2CO3), Phosphoric acid (H3PO4), etc. On the basis of Concentration On the basis of Concentration, acids are categorized into two categories Strong Acids Weak Acids Strong Acids A strong Acid is an acid which completely ionizes in water and produces (H+). For example, Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3) etc. Weak Acids A weak acid is partially ionized in water, creating a tiny amount of hydrogen ions (H+). For example, Acetic acid (CH3COOH), Carbonic acid (H2CO3) etc. 6.5 Bases A compound that can neutralize an acid and produces salt and hydroxide((OH–). Bases turn red litmus paper blue while the blue litmus paper stays blue. They taste bitter and also feel soapy to the touch. Aqueous solutions of the bases can conduct electricity. Some common examples are; copper oxide, sodium hydroxide sodium bicarbonate (used in cooking, and household bleach). 6.6 Chemical Properties of Base few of the following chemical properties of bases are, 47  6.6.1 Reaction of Base with Metals: When bases react with metal, salt and hydrogen gas is produced. Alkali + Metal → Salt + Hydrogen When sodium hydroxide interacts with aluminium metal, sodium aluminate and hydrogen gas are generated. 2NaOH + 2Al + 2H2O → 2NaAlO2 + 2H2 6.6.2 Reaction of Non-Metallic Oxides with Base: Salt and water are formed when non -metallic oxides react with a base. Non-metallic oxide + Base → Salt + Water When calcium hydroxide reacts with carbon dioxide calcium carbonate is formed along with water. Ca(OH)2 + CO2 → CaCO3 + H2O 6.6.3 Reaction of Bases with Ammonium Salts: Ammonia is produced when alkalis react with ammonium salts. Alkali + Ammonium salt → Salt + Water + Ammonia When calcium hydroxide reacts with ammonium chloride, calcium chloride, water, and ammonia are produced. Ca(OH)2 + NH4Cl → CaCl2 + H2O + NH3 6.7 Types of Bases Acidity, concentration, and degree of ionization are three variables that can be used to classify bases. Types of Bases Based on Acidity The number of hydroxyl ions presents determines acidity in bases. Based on acidity, bases are classified into three categories: Monoacidic Diacidic Triacidic 48  6.7.1 Monoacidic bases are those that contain only one hydroxyl ion and interact with only one hydrogen ion. Mono-acidic bases include NaOH, KOH, and others. 6.7.2 Diacidic base is a base with two hydroxyl ions that interact with two hydrogen ions. Ca (OH)2, Mg (OH)2, and other di-acidic bases are examples. 6.7.3 Triacidic a type of base that comprises three hydroxyl ions and three hydrogen ions. And include Al (OH)3, Fe (OH)2, and others. 6.8 Types of Bases Based on their Concentration Based on their concentration in an aqueous solution, bases are divided into two categories: Concentrated Diluted Concentrated: The concentration of base is higher in the solution in their aqueous solution. Concentrated NaOH solution, for example. Diluted: These types of bases have a lower concentration of base in their aqueous solution. For instance, dilute NaOH, dilute KOH, and so on. 6.9 Types of Bases Based on their Degree of Ionization When bases are dissolved in water, it produces a certain quantity of hydroxyl ions. The degree of ionization distinguishes two types of bases. Strong Base Weak Base Strong Base: A strong base is one that dissociates entirely or to a large extent in water. For example, NaOH, KOH, and strong bases. Weak Base: A weak base is one that does not dissolve entirely or only dissociates to a very little level. For example, NH4OH, and others are weak bases. 6.10 Alkali Bases that are easily dissolved in water are called Alkali, in other words, water soluble bases are called Alkali. For example, NaOH is an alkali as it dissolves in water forming Na+ and OH– ions. 49 The three most important modern concepts of acids and bases are: Arrhenius Concept According to this concept, Substances which produce H+ ions when dissolved in water are called acids while those which ionize in water to produce OH – ions are called bases. HA → H+ + A– (Acid) BOH → B+ + OH– (Base) Arrhenius proposed that reactions are acidic if they dissociate in aqueous solution to form hydrogen ions (H+) and bases if they form hydroxide (OH–) ions in aqueous solution. Limitations of Arrhenius Concept The presence of water is absolutely necessary for acids and bases. Dry HCl can’t act as an acid. HCl acts as an acid in water only and not in any other solvent. The concept does not explain the acidic and basic character of substances in non-aqueous solvents. The neutralization process is only possible for reactions which can occur in aqueous solutions, although reactions involving salt formation can occur in the absence of a solvent. Bronsted-Lowry Concept.According to them, an acid is defined as any hydrogen-containing material (molecule, anion or cation) which can donate a proton to other substance and a base is any substance (molecule, cation or anion) that can accept a proton from any other substance. Therefore, acids are proton donors whereas bases are proton acceptors. 50 Conjugate Acid-Base Pairs Consider a reaction Acid1 + Base2 → Acid2 + Base1 H2O + HCl ⇔ H3O+ + Cl– In this reaction, HCl donates a proton to H2O and is, therefore an acid. Water, on the other hand, accepts a proton from HCl, and is, therefore, a base. In the reverse reaction, the H3O+ ions donate a proton to Cl– ion, hence H3O+, an ion is an acid. Cl– ion, because it accepts a proton from H3O+ ion, is a base. Acid-base pairs in which the members of the reaction can be formed from each other by gaining or losing protons are called conjugate acid-base pairs. Limitations of Bronsted Lowry Concept Bronsted Lowry could not explain the reaction occurring in non-protonic solvent like COCl3, SO2, N2O4, etc. It cannot explain the reactions between acidic oxides and the basic oxides which can easily take place in the absence of solvent as well e.g. (No proton transfer) Substances like BF3, AlCl3 , etc, do not contain hydrogen which means they can’t donate a proton, still they behave as acids. 51 Lewis Concept According to this theory bases donate pairs of electrons and acids accept pairs of electrons. Thus, it can be said that a Lewis acid is electron-pair acceptor. Oxidation-reduction reactions take place on a transfer of electrons from one atom to another, with a net change in the oxidation number of one or more atoms. There is no change in the oxidation numbers of any atoms. Either an electron is transferred from one atom to another, or the atoms come together to share a pair of electrons. Al (OH)3 + 3H+ → Al3+ + 3H2O (Aluminium hydroxide is acting as a base) Al (OH)3 + OH– → Al (OH)4- (Aluminium hydroxide is acting as an acid) When Aluminium hydroxide accepts protons, it acts as a base. When it accepts electrons, it acts as an acid. This theory also explains why non-metal oxides such as carbon dioxide dissolve in H2O to form acids, such as carbonic acid H2CO3. CO2(g) + H2O(l) → H2CO3(aq) Limitations of Lewis Concept Lewis's concept gave a generalized idea including all coordination reactions and compounds. This is not always true. An idea about the relative strength of acids and bases is not provided by Lewis's concept. Lewis's concept is not in line with the acid-base reaction concentration. 6.11 Salts When an acid and a base react, they generate salts, which are ionic substances. Salts do not have an electrical charge. Apart from sodium chloride, other common salts are sodium nitrate, barium sulfate etc. Sodium chloride or common salt is a product of the reaction between the hydrochloric acid (acid) and sodium hydroxide (base). Solid sodium chloride is made of a cluster of positively charged sodium ions and negatively charged chloride ions held together by electrostatic forces. Physical Properties of Salt some of physical properties of salts are, The bulk of the salts are crystalline. Salts that are transparent or opaque are available. The salts are soluble in water. Salt solutions, in their molten state, also transmit electricity. The taste of salts can be salty, sour, sweet, bitter, or savoury. There is no odour to neutral salts. 52  Because it contains ions, salt water is an excellent conductor of electricity. Electrostatic attraction holds the ions together, and a chemical bond is established between them. Types of Salts Some of the important categories of salts are given below Acidic Salts Basic or Alkali Salts Neutral Saltss 6.11.1 Acidic Salt A class of salts that produce an acidic solution after being dissolved in a solvent. Its formation as a substance has a greater electrical conductivity than that of the pure solvent. An acidic solution formed by acid salt is made during partial neutralization of diprotic or polyprotic acids. A partial neutralisation of a diprotic or polyprotic acid produces an acidic salt. These salts contain H+ cations or strong cations in their aqueous solution. The ionizable H+ makes up the majority of the ions. Some examples of acidic salts are NaHSO4, KH2PO4 etc. These salts are formed by the neutralization of strong acids and weak bases. Ammonium Chloride Ammonium chloride is formed when hydrochloric acid (a strong acid) interacts with ammonium hydroxide (a weak base). NH4OH + HCl → NH4Cl + H2O Ammonium Sulphate Ammonium sulphate is formed when ammonium hydroxide (a weak base) reacts with sulphuric acid (a strong acid). 2NH4OH + H2SO4 → (NH4)2SO4 + 2H2O 6.11.2 Basic or Alkali Salt A basic salt is formed when a strong base reacts with a weak acid to partially neutralise it. When they are hydrolyzed, they decompose into a basic solution. This is because when a basic salt is hydrolyzed, it produces the conjugate base of a weak acid in the solution. e.g. Sodium Carbonate (Na2CO3), Sodium Acetate (CH3COONa). Sodium Carbonate Sodium carbonate is formed when sodium hydroxide (a strong base) reacts with carbonic acid (a weak acid) H2CO3 + 2 53 NaOH → Na2CO3 + H2O Sodium Acetate Sodium acetate is formed when a strongly basic, sodium hydroxide (a strong base), reacts with acetic acid (a weak acid) CH3COOH + NaOH → CH3COONa + H2O 6.11.3 Neutral Salts Salts generated by the reaction of a strong acid with a strong base are neutral in nature. The pH of these salts is 7, which is considered neutral. Potassium Chloride, Sodium Chloride, and others are examples of neutral salts. Sodium Chloride is formed when hydrochloric acid (a strong acid) mixes with sodium hydroxide (a strong base). NaOH + HCl → NaCl + H2O Salts can also be categorised into other categories which include, Mixed Salts Double Salt 6.11.4 Double Salt Salts with more than one cation or anion are known as double salts. They’re created by mixing two different salts that crystallised in the same ionic lattice. e.g. Potassium Sodium Tartrate (KNaC4H4O6.4H2O) also known as Rochelle salt. 6.11.5 Mixed Salts; Salts which are produced by mixing two salts, which generally share a common cation or anion, are called mixed salts. e..g. Bleaching Powder CaOCl2. 6.11.6 Formation of Acidic, Basic, and Neutral Salts When a strong acid reacts with a weak base, the base is unable to completely neutralize the acid. As a result acidic salt is formed. When a strong base is combined with a weak acid, the acid is unable to completely neutralize it. As a result a simple salt is formed. When an equal-strength acid and base react, they totally neutralise each other. A neutral salt is formed as a result of this process. 6.12 Strength of Acids and Base The strength of an acid or a base is measured by the amount of H+ ions or OH– ions present in the aqueous solutions. Strong acids have a higher concentration 54 of H+ ions per unit volume in their aqueous solution whereas weaker acids or bases have a lower concentration of H+ ions or OH- per unit volume in their aqueous solutions. An acid’s strength is affected by the electronegativity of the conjugate base and the polarity of the acidic hydrogen. Strength refers to how readily the hydrogen cation (H+) disassociates from the anion. Strong acids and bases dissociate entirely in aqueous solutions, whereas weak acids and bases dissociate partially into their conjugate ions. The strength of an acid or a base is measured by the amount of H+ ions or OH– ions present in their aqueous solution. The strength of Aicds and Bases can easily be measured using a pH scale.. It is calculated using the formula, pH = -log[H+] For an acid, pH ranges from 0 to 7 whereas for a base it ranges between 7 and 14. The lower the pH higher is the strength of the acid and the higher the pH higher the strength of the base. 6.13 Indicator Indicators are chemical compounds which help to indicate the presence of an acid or a base in a chemical reaction. They possess different colors in acidic solutions and different colors in basic solutions. Indicators are made naturally by plants and animals or artificially. The image shows a litmus test of ids and bases. 55 Types of Indicators Various types of indicators used are mentioned below, Natural Indicators: are derived from plants, animals or any living organism, examples, Red Cabbage, Litmus paper. Synthetic Indicators: these are made artificially in laboratories and factories are synthetic indicators, examples, are Phenopthelien, Methyl orange. Olfactory Indicators: Substances that have different smells in an acidic or basic medium are Olfactory Indicators, example onions, olives and others. 6.14 Titration In titration, a solution of a known concentration, called a standard solution, is used to determine the concentration of another solution. For acid-base titrations, a standardized solution of base is slowly added to an acid of unknown concentration or the acid is added to the base. The acid-base reaction is a neutralization reaction, which forms a salt and water. When the moles of hydrogen ions in the acid are equal to the moles of hydroxyl ions added from the base, the solution reaches neutral pH. To perform an acid-base titration, the standardized base is slowly added to a stirring flask of the unknown acid using a burette, which enables the measurement of volume and the dropwise addition of base. The pH of the solution is closely monitored throughout the titration using a pH indicator added to the acid. Typically, phenolphthalein is used as the solution remains colorless until it becomes basic, turning a light pink. As the equivalence point is reached, (when the moles of hydrogen ions equal the moles of hydroxyl ions), the pH indicator temporarily changes color due to an excess of hydroxyl ions. When the flask is swirled, the pH indicator’s acidic color returns. The titration is complete and has reached its endpoint when a tiny excess of hydroxyl ions changes the indicator permanently to its basic color. The titration curve is a plot of the pH of a solution versus the volume of standardized base added. The equivalence point is located at the inflection point of the curve, and it is calculated as the second derivative of the titration curve. If an acid is polyprotic, it will have multiple equivalence points, one for each hydrogen ion dissociation. The pH at the halfway point to the equivalence point for monoprotic acids, or between equivalence points in the case of polyprotic acids, is equal to the pKa of the acid. 56  6.15 pH pH is a measure of the amount of hydrogen ions in a solution and the degree of the acidity of the solution. The pH range runs from 0 to 14; aqueous solutions with a pH below 7 are acidic, and aqueous solutions with a pH above 7 are alkaline or basic. Solutions at pH 7 are considered neutral. The pH of a solution is equal to the negative log base ten of the concentration of hydrogen ions in solution. Water interacts strongly with the hydrogen ion because its strong positive charge attracts the negative pole of surrounding water molecules. In fact, they interact so strongly that they form a covalent bond and the H3O+ cation, called hydronium. The above equation is rewritten to reflect this. For simplicity, we’ll refer to the concentration of hydrogen ions instead of hydronium ions when discussing pH. The lower the pH value of a solution, the more hydrogen ions that are present, and the more acidic the solution. For example, the pH of 1 mM of sulfuric acid is 2.75, whereas the pH of 1 mM of hydrochloric is 3.01. The concentration of hydrogen ions in the sulfuric acid solution is calculated as 1 × 102.75, whereas the concentration of hydrogen ions in the hydrochloric acid solution 57 is 1 × 10-3.01. Thus, there are more hydrogen ions present in sulfuric acid, and it is more acidic. A-. Higher Ka values represent stronger acids, whereas smaller Ka values represent weaker acids. Ka is reported in the form of pKa, which