2. GENERAL CHEMISTRY (CHEM.1012) By: Adisu G. MSc. (In Inorganic Chemistry) Madda Walabu University November, 2022 GC Bale Robe, Ethiopia “Excellence Through Diversity!” College of Natural and Computational Science, Department of Chemistry 1

3. INTRODUCTION  Practical Art and Crafts (…to 600 B.C.E)  Greek Period (600 to 300 B.C.E).  Alchemy (300 B.C.E to 1650 C.E)  Phlogiston (1650 to 1790)  Modern Chemistry (1790…) Practical Art and Crafts Variation found in almost all cultures Practical uses of chemistry in “everyday life” pottery, baking, brewing, Medicines, extracting metals from ores Greek Period Philosophical approach Attempt to understand principles that would explain the universe Not supported by empirical evidence Not supported by experimental data 2 History of Chemistry

4. Alchemy Combined Greek philosophical traditional with practical Egyptian craft tradition Mixed astrology and mysticism and Greek philosophical traditional with practical knowledge Main focus of Alchemists Search for philosophers stone (elixir of life) Confer immorality Transmute base metals into gold Later emphasis on latrochemistry (medicinal chemistry) Phlogiston Era Stahl thought that there was a” fire principle” called phlogiston in anything that would burn Phlogiston would be released when it burned Wood Ash + Phlogiston (into air) 3 Cont…

5. Modern Chemistry Robert Boyle- the skeptical chemist 1661 relied on experimental evidence Lost of information developed during the intervening years Antoine Lavoisier (1743-1794)  Carefully weighed materials before and after burning (closed container) Elementary treatise on chemistry ….1789 chemistry described in ‘ modern” terms Law of conservation of matter  Matter cannot be created or destroyed by ordinary chemical means  Alchemy was based more on experimentation and had little basis in science.  Chemistry utilizes both experimentation and scientific practices.  Modern chemistry basically relies on scientific theories and experimental results, but the alchemy was a blend of myths, religion, magic, astrology, philosophy, and spirituality 4 Cont…

6. CHAPTER ONE ESSENTIAL IDEAS IN CHEMISTRY science that deals with the properties, composition, and structure that transformations they undergo, and the energy that is released or absorbed during these processes.  What is Chemistry? Is science that deals with the properties, composition, and structure that transformations they undergo, and the energy that is released or absorbed during these processes.  In most general term, Chemistry is the study of chemicals the study of chemicals the science that explains how matter changes. the science that explains how matter changes. is the science describing matter and its transformations and is central to virtually all areas of modern science and technology.  Why should you/we study chemistry? is the science describing matter and its transformations and is central to virtually all areas of modern science and technology. 5

7.  Throughout human history, people have tried to convert matter into more useful forms.  Our Stone Age ancestors chipped pieces of flint into useful tools and carved wood into statues and toys.  These endeavors involved changing the shape of a substance without changing the substance itself.  But as our knowledge increased, humans began to change the composition of the substances as well clay was converted into pottery, hides were cured to make garments, hides were cured to make garments, copper ores were transformed into copper tools and weapons, and copper ores were transformed into copper tools and weapons, and grain was made into bread. grain was made into bread. 1.1. Chemistry in Context 6

8. Humans began to practice chemistry to:  Control fire and use it to cook, make pottery, and smelt metals. bronze  Combination of metals to form alloys; eg: copper and tin were mixed to make bronze and more elaborate smelting techniques produced iron.  Extracted Alkalis from ashes  Prepared soap by combining alkalis with fats (Na, K,… with fatty acids).  Began to separate and use specific components of matter. Eg: variety of drugs such as aloe, myrrh, and opium were isolated from plants and Dyes (indigo and Tyrian purple) were extracted from plant and animal.  Alcohols can be synthesized by the techniques of fermentation, traditional way of distillation,…  From alchemy to modern chemistry; chemistry continues to deepen our understanding and improve our ability to couple and control the behavior of matter and universality of chemistry in daily life. Cont… 7

9. 1.1.1. Chemistry as the Central Science  Chemistry is sometimes referred to as “the central science” due to its interconnectedness with a vast array of other STEM disciplines (STEM stands for areas of study in the Science, Technology, Engineering, and Mathematics fields).  Chemistry is one of the most important of all science for human welfare, b/c knowledge of the principles of chemistry can facilitate understanding of other sciences  Chemistry and the language of chemists play vital roles in biology, medicine, materials science, nanotechnology, forensics, Chemical engineering, environmental science, and many other fields of study.  Environmental science, geology, oceanography, and atmospheric science incorporate many chemical ideas to help us better understand and protect our physical world.  Chemical ideas are used to help understand the universe in astronomy and cosmology 8

10. Figure 1.1: Knowledge of chemistry is central to understanding a wide range of scientific disciplines. This diagram shows just some of the interrelationships between chemistry and other fields. 9 Cont…

11.  A systematic approach to research, to investigate, to cast about, to beat about, to re-explore in order to solve problem.  The scientific method is a process that scientists use to make observations in nature, gather data, and explain natural phenomena.  Chemistry is a science based on “observation and experimentation”: involves attempting to answer questions and explain observations in terms of the laws and theories of chemistry, using procedures that are accepted by the scientific community. involves attempting to answer questions and explain observations in terms of the laws and theories of chemistry, using procedures that are accepted by the scientific community. no single route to answering a question or explaining an observation, but there is an aspect common to every approach. no single route to answering a question or explaining an observation, but there is an aspect common to every approach. 1.1.2. The Scientific Method Cont… 10

12. Cont… Sequential steps in the scientific method Step-1: Observations  The first step in the scientific method is to make observations about nature and ask questions about what you observe.  When an observation always seems to be true, it may be stated as a law that predicts that behavior and is often measurable.  However, a law does not explain that observation.  For example, we can use the Law of Gravity to predict that if we drop our chemistry book it would fall on the table or the floor, but this law does not explain why our book falls Step.2: Hypothesis  A tentative explanation of observations that acts as a guide for gathering and checking information.  A scientist forms, which gives a possible explanation of an observation or a law. 11

13.  It can be tested by experiments, calculation, or comparison with the experiments of others and then refine it as needed. Step-3: Experiments To determine if a hypothesis is true or false, experiments are done to find a relationship between the hypothesis and the observations. The results of the experiments may confirm the hypothesis. Cont… 12

14.  However, if the experiments do not confirm the hypothesis, it is modified or discarded.  Then new experiments will be designed to test the hypothesis. Step-4: Conclusion/Theory  are well-substantiated, comprehensive, testable explanations of particular aspects of nature.  When the results of the experiments are analyzed, a conclusion is made as to whether the hypothesis is true or false.  When experiments give consistent results, the hypothesis may be stated to be true.  Even then, the hypothesis continues to be tested and, based on new experimental results, may need to be modified or replaced.  If many additional experiments by a group of scientists continue to support the hypothesis, it may become a scientific theory, which gives an explanation for the initial observations. Cont… 13

15. 1.1.3. The Domains of Chemistry Cont…  Chemists study and describe the behavior of matter and energy in three different domains.  These domains provide different ways of considering and describing chemical behavior:  macroscopic - is realm of everyday things that are large enough to be sensed directly by human sight or touch. Eg Density, Solubility, and Flammability  Microscopic- is often visited in the imagination. Some are visible through standard optical (sophisticated) microscopes, capable of imaging even smaller entities such as biological cell, molecules and atoms, ions and electrons, protons and neutrons, chemical bonds, each of which is far too small to see symbolic 14

16.  Symbolic- contains the specialized language used to represent components of the macroscopic and microscopic domains. These symbols play an important role in chemistry, b/c they help to interpret the behavior of the macroscopic domain in terms of the components of the microscopic domain. Eg- chemical formula, Chemical equation, periodic table symbols, …. 15 Cont…

17. 1.2. State and Classification of Matter 1.2.1. State of Matter Matter is anything that occupies space and has mass, and chemistry is the study of matter and the changes it undergoes. solid, liquid, and gas plasma  All matter, at least in principle, can exist in three/four states: solid, liquid, and gas on earth and plasma  Solids- are rigid objects  Liquids- are less rigid than solids  Gases- takes both the shape and volume of its container, they can expand indefinitely.  Plasma- is a gaseous state of matter that contains appreciable numbers of electrically charged particles 16

18. 17 STATES OF MATTER SOLID LIQUID GAS PLASMA Tightly packed, in a regular pattern Vibrate, but do not move from place to place Close together with no regular arrangement. Vibrate, move about, and slide past each other Well separated with no regular arrangement. Vibrate and move freely at high speeds Has no definite volume or shape and is composed of electrical charged particles

19. Classification of Matter can be classified into several categories. Two broad categories are mixtures and pure substances. Matter- can be classified into several categories. Two broad categories are mixtures and pure substances. “there is no detectable change in the total quantity of matter present when matter converts from one type to another (a chemical change) or changes among solid, liquid, or gaseous states (a physical change)”. The law of conservation of matter summarizes many scientific observations about matter: It states that “there is no detectable change in the total quantity of matter present when matter converts from one type to another (a chemical change) or changes among solid, liquid, or gaseous states (a physical change)”. Cont… 18

20. 1.3. Physical and Chemical Properties properties  The characteristics that distinguish one substance from another are called properties.  Substances are identified by their properties as well as by their composition.  Properties of a substance may be quantitative (measured and expressed with a number) or qualitative (not requiring explicit measurement).  Physical Properties: can be observed without changing the identity and composition of the substance. Familiar physical properties are color, odor, density, melting point, boiling point, and hardness. Familiar physical properties are color, odor, density, melting point, boiling point, and hardness.  Chemical properties: describe the way a substance may change, or react, to form new substances. A common chemical property are flammability, the ability of a substance to burn in the presence of oxygen. A common chemical property are flammability, the ability of a substance to burn in the presence of oxygen. 19

21. The changes substance undergo are either physical or chemical:  During a physical change, a substance changes its physical appearance but not its composition. (same substance before and after the change). Melting is a physical change: one in which the state of matter changes, but the identity of the matter does not change. Melting is a physical change: one in which the state of matter changes, but the identity of the matter does not change. We can recover the original ice by cooling the water until it freezes. We can recover the original ice by cooling the water until it freezes.  In a chemical change ( chemical reaction), a substance is transformed into a chemically different substance. Hydrogen gas burns in oxygen gas to form water  The statement “Hydrogen gas burns in oxygen gas to form water” describes a chemical property of hydrogen, b/c to observe this property we must carry out a chemical change—burning in oxygen (combustion).  After a chemical change, the original substance (hydrogen gas in this case) will no longer exist.  We cannot recover the hydrogen gas from the water by means of a physical process, such as boiling or freezing Cont… 20

22.  All measurable properties of matter fall into two categories: extensive properties and intensive properties.  Extensive property: Property of matter that is quantity dependent Eg : Mass, length, and volume Values of the same extensive property can be added together  Intensive property: Property of matter that is independent quantity Eg. Color, odor, hardness. conductivity, density, melting, freezing, boiling point, Temperature etc…  Unlike mass and volume, temperature and other intensive properties such as melting point, boiling point, and density are not additive  The ratio of two extensive properties is an intensive property: D = m/v 1.4. Extensive and intensive property 21

23. 1.5. Measurements and Units  To understand any phenomenon in chemistry we have to perform experiments Experiments require measurements, and we measure several physical properties like length, mass, time, temperature, pressure etc. Experimental verification of laws & theories also needs measurement of physical properties 22

24. Cont…  A scientists use a variety of devices to measure the properties of matter.  Measurements provide information's about the hypotheses, theories, and laws: describing the behavior of matter and energy in both the macroscopic and microscopic domains of chemistry. number, unit and the uncertainty Every measurement provides three kinds of information: number, unit and the uncertainty.  Qualitative properties : measurements are words, such as heavy or hot  Quantitative properties : measurements involves number (quantities), and depends on : the reliability of the measuring instrument and the care with which it read 23

25. Cont… SI Base Units  The revised metric system is called the International System of Units (abbreviated SI, from the French (Système Internationale d’Unités).  There are seven SI base units.  All other units of measurement can be derived from these base units. The standard used for the measurement of a physical quantity is called a unit. Table 1. Examples of SI Base units 24

26. SI Derived Units Cont…  Other quantities, called derived quantities, are defined in terms of the seven base quantities via a system of quantity equations.  The SI derived units for these derived quantities are obtained from these equations and the seven SI base units.  Examples of such SI derived units are given in Table 2, Table 2. Examples of SI derived units SI derived unit Derived quantityNameSymbol areasquare meterm2m2 volumecubic meterm3m3 Density kilogram per cubic meter kg/m 3  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 25

27.  Density is the ratio of mass to volume. Oil floats on water, because in addition to not mixing with water, oil has a lower density than water. That is, given equal volumes of the two liquids, the oil will have a smaller mass than the water. is calculated using the following equation:  The SI unit for density is (kg/m 3 ).  For many situations, we often use (g/cm 3 ) for the densities of solids and liquids, and (g/L) for gases. Example a.To three decimal places, what is the volume of a cube (cm 3 ) with an edge length of 0.843 cm? Soln:______ b. If the cube in part (a) is copper and has a mass of 5.34 g, what is the density of copper to two decimal places? Soln:_____ 26 Cont…

28. 1.6. Uncertainty in Measurement  Chemistry makes use of two types of numbers: exact and inexact complete certainty  Exact numbers- is a value that is known with complete certainty. include numbers with defined values, such as 1 inch = 2.54 cm, 1 kg = 1000 g, and 1 dozen = 12 objects. (has zero uncertainity and cannot be simplified/reduced.) Exact numbers also include those that are obtained by counting.  Inexact numbers- has a value that has a degree of uncertainty associated with it. are generated anytime when a measurement is made. are generated anytime when a measurement is made. measured by any method other than counting. measured by any method other than counting. Measured values, estimates, rounded numbers and some unit conversions are an examples. Measured values, estimates, rounded numbers and some unit conversions are an examples.  The uncertainty of a measurement tells us something about its quality. is the doubt that exists about the result of any measurement. 27

29. 1.6.1. Significant Figures in Measurement Cont…  Significant figures express the uncertainty of a measurement or number.  All measurements have some degree of uncertainty in their value An inexact number must be reported in such a way as to indicate the uncertainty in its value. This is done using significant figures. Significant figures are the meaningful digits in a reported Number.  In a measured number, the significant figures (SFs) are all the digits including the estimated digit. significant figures  Non-zero numbers are always counted as significant figures. zero may or may not be a significant figure position in a number However, a zero may or may not be a significant figure depending on its position in a number. We will use the terms “leading,” “trailing,” and “captive” for the zeros and will consider how to deal with them. 28

30. Cont… 29

31. Cont… RULES FOR SIGNIFICANT FIGURES ARE 1. All non-zero numbers ARE significant. The number 33.2 has THREE significant figures because all of the digits present are non-zero. ARE 2. Zeros between two non-zero digits ARE significant. 2051 has FOUR significant figures. The zero is between a 2 and a 5. NOT significant 3. Leading zeros are NOT significant. They're nothing more than place holders. 0.0032 also has TWO significant figures. All of the zeros are leading. ARE significant 4. Trailing zeros to the right of the decimal ARE significant. There are FOUR significant figures in 92.00, 92.00 is different from 92. a scientist who measures 92.00 mL knows his value to the nearest 1/100th mL; meanwhile his colleague who measured 92 mL only knows his value to the nearest 1 mL. It's important to understand that "zero" does not mean "nothing." Zero denotes actual information, just like any other number. ARE significant NOT significant 5. Trailing zeros in a whole number with the decimal shown ARE significant. Placing a decimal at the end of a number is usually not done. By convention, however, this decimal indicates a significant zero. For example, "540." indicates that the trailing zero significant; there are THREE significant figures in this value; but with no decimal "540" is NOT significant 30

32. Example: Identify the significant zeros in each of the following measured numbers: a. 0.000 250 m b. 70.040 g c. 1 020 000 L Solution a. The zero in the last decimal place following the 5 is significant. The zeros preceding the 2 are not significant. b. Zeros between nonzero digits or at the end of decimal numbers are significant. All the zeros in 70.040 g are significant. c. Zeros between nonzero digits are significant. Zeros at the end of a large number with no decimal point are placeholders but not significant. The zero between 1 and 2 is significant, but the four zeros following the 2 are not significant. 6. Exact numbers have an INFINITE number of significant figures. Example, 1 meter = 1.00 meters = 1.0000 meters =1.000000 meters, etc. 7. For a number in scientific notation: N x 10 x, all digits comprising N ARE significant by the first 6 rules; "10" and "x" are NOT significant. 5.02 x 10 4 has THREE significant figures: “10” and "4" are not significant. Cont… 31

33. Cont… 1.6.2. Significant Figures in Calculations Because we often use one or more measured numbers to calculate a desired result, a second set of guidelines specifies how to handle significant figures in calculations. 1.Rules for Rounding Off 1. If the first digit to be dropped is 4 or less, then it and all following digits are simply dropped from the number. 2. If the first digit to be dropped is 5 or greater, then the last retained digit of the number is increased by 1. 32

34. 2. In addition and subtraction, the answer cannot have more digits to the right of the decimal point than the original number with the smallest number of digits to the right of the decimal point. For example: 102.50 ← two digits after the decimal point + 0.231 ← three digits after the decimal point 102.731 ← round to 102.73 143.29 ← two digits after the decimal point − 20.1 ← one digit after the decimal point 123.19 ← round to 123.2 Cont… 33

35. 3.In multiplication and division- the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of significant figures. The following examples illustrate this rule: 1.4 × 8.011 = 11.2154 ← round to 11 (limited by 1.4 to two significant figures) 11.57 305.88 = 0.037825290964 ← round to 0.03783 (limited by 11.57 to four significant figures) Cont… 34

36. 4. Exact numbers- can be considered to have an infinite number of significant figures and do not limit the number of significant figures in a calculated result. For example, a penny minted after 1982 has a mass of 2.5 g. If we have three such pennies, the total mass is 3 × 2.5 g = 7.5 g The answer should not be rounded to one significant figure because 3 is an exact number. Cont… 35

37. 4. In calculations with multiple steps, rounding the result of each step can result in “rounding error.” Consider the following two-step calculation: First step: A × B = C Second step: C × D = E Suppose that A = 3.66, B = 8.45, and D = 2.11. The value of E depends on whether we round the value of C prior to using it in the second step of the calculation (Method 1) or not (Method 2). Method 1 Method 2 C = 3.66 × 8.45 = 30.9 C = 3.66 × 8.45 = 30.93 E = 30.9 × 2.11 = 65.2 E = 30.93 × 2.11 = 65.3 In general, it is best to retain at least one extra digit until the end of a multistep calculation, as shown by Method 2, to minimize rounding error. Cont… 36

38. 1.6.3. Accuracy and Precision  Accuracy and precision are two important factors to consider when taking data measurements.  Both accuracy and precision reflect how close a measurement is to an actual value.  Although the difference between the two terms may be subtle, it is important.  Accuracy- tells us how close a measurement is to the true value.  Precision- tells us how close reproducible measurements of the same thing are to one another; even if they are far from the accepted value ( see figure below).  Measurements are said to be precise if they yield very similar results when repeated in the same manner.  A measurement is considered accurate if it yields a result that is very close to the true or accepted value. Cont… 37

39. 1.7. Conversion Factors and Dimensional Analysis  A ratio of two equivalent quantities expressed with different measurement units can be used as a unit conversion factor  The linear equation relating Celsius and Fahrenheit temperatures is easily derived from the two temperatures used to define each scale. 38

40. CHAPTER 2 ATOMS, MOLECULES AND IONS 39

41. INTRODUCTION Atoms are the smallest unit of matter that can’t be broken down chemically Molecules are groups of two or more atoms that are chemically bonded. Ions are atoms or molecules that have gained or lost one or more of their valence electrons and therefore have a net positive or negative charge. 40  An atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. Atom = (neutron+ proton) + electron.

42. 41 History of the Atom 1. Democritus (400 BC) Believed universe made of invisible units called atoms Named them Atoms Greek word atomos, means something invisible and cannot be cut or divided

43. 42 Aristotle was part of the generation that succeeded Democritus. He did not believe in atomos. Aristotle thought that all matter was continuous. That is, if one proceeded on breaking down a substance, it would be impossible to reach to the last indivisible particle. In other words, it would continue to divide infinitely. His opinion was accepted for nearly 200 years Aristotle  The early concept of atoms was simply a result of thinking and reasoning on the part of the philosophers, instead of experimental observations.  In 1803, however, John Dalton proposed a completely different theory of matter.  His theory was based on scientific experimental observations and logical laws

44. Cont… 43 2. Dalton’s Atomic Theory (1807) John Dalton (1766 – 1844) Wrote the first atomic theory 1. All elements are composed of tiny indivisible particles called atoms 2.Atoms of the same element are identical. 3. Atoms of any one element are different from those of any other element. 3.Atoms of different elements combine in simple whole-number ratios to form chemical compounds 4.Atoms cannot be created or destroyed. 5. In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

45. Cont… 44 during the 1900s evidence was discovered regarding charges: – atoms have positive (Rutherford’s contribution) and negative (Thomson’s contribution) parts – charges interact: as a result, revisions to Dalton’s model had to be made New Evidence

46. Cont… 45 3. J.J. Thomson (In 1897) used a cathode ray tube to deduce the presence of a negatively charged particle: the electron cathode ray: beam of negative particles

47. Cont… 46 Conclusions from the Study of the Electron: A. Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. B. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons C. Electrons have so little mass that atoms must contain other particles that account for most of the mass

48. Cont… 47  Nuclear Model discovered a huge flaw in the previous concept of the atom during his now famous gold foil experiment Discovered the Nucleus and the Positive Proton Surmised atoms are made of mostly empty space Didn’t know about the Neutrons Famous Gold Foil Experiment 4. Ernest Rutherford (1911)

49. Cont… 48 Gold Foil Experiment Particles shot through thin sheet of gold Most shots went straight through A small amount were deflected Hence… The atoms must be made of mostly empty space with a small dense nucleus #1 The nucleus is small #2 The nucleus is dense #3 The nucleus is positively charged Conclusions:

50. Cont… 49 5. James Chadwick (1932) Chadwick named these particles neutrons. Chadwick named these particles neutrons. He bombarded a thin foil of beryllium with α-particles of a radioactive substance. He then observed that highly penetrating rays, consisting of electrically neutral particles of a mass approximately equal to that of the proton, were produced. These neutral particles are called neutrons

51. Cont… 50 The Modern Atomic Theory can be summarized as follows: 1.Atoms are the smallest particles of all elements that can take part in a chemical reaction. 2.An atom is divisible. It can be subdivided into electrons, protons, and neutrons. 3.An atom is also indestructible i.e., atoms can neither be created nor destroyed during ordinary chemical reactions. Atoms of the same element may not be identical in mass because of the existence of isotopes. 4.Atoms of the same elements have identical chemical properties. 5.Atoms of different elements have different chemical properties. 6.Atoms of two or more elements combine in simple whole-number ratios to form compounds.

52. 51 2.1. Atomic structure and symbolism  Series of investigations that began in the 1850s and extended into the twentieth century clearly demonstrated that atoms actually possess internal structure: Internal structure of atoms are made up of even smaller particles, which are called subatomic particles. This research led to the discovery of three particles such as electrons, protons, and neutrons. An atom consists of a central nucleus that is usually surrounded by one or more electrons. Each electron is negatively charged.

53. 52  The total number of protons and neutrons in an atom is called its mass number (A)  Symbol : Where, E= Element Symbol Z= atomic number= # of proton = # of electron (unless it is charged) A= atomic mass Example Z=11, A=23 Cont…

54. 2.1.1. Chemical Symbols and Isotopes 53  A chemical symbol: is an abbreviation that we use to indicate an element or an atom of an element. For example, the symbol for mercury is Hg.  We use the same symbol to indicate one atom of mercury (microscopic domain) or to label a container of many atoms of the element mercury (macroscopic domain).  Some symbols are derived from the common name of the element; others are abbreviations of the name in another language.  Most symbols have one or two letters, but three-letter symbols have been used to describe some elements that have atomic numbers greater than 112.  To avoid confusion with other notations, only the first letter of a symbol is capitalized.  For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen(O).

55. 54

56. 55 same atomic number different numbers of neutrons  Isotopes are atoms of the same element that have the same atomic number but different numbers of neutrons. Cont…

57. 56 Frederick Soddy Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials. Cont…

58. 57 2.1.2. Atomic mass unit and average atomic mass  According to international agreement, one atomic mass unit (amu) is defined as a mass exactly equal to one-twelfth the mass of one carbon-12 atom.  (Carbon-12 is the carbon isotope that has six protons and six neutrons), Setting the atomic mass of carbon-12 at 12 amu provides the standard for measuring the atomic mass of the other elements. hydrogen atom (1H) is only 8.3985 percent  For example, experiments have shown that a hydrogen atom (1H) is only 8.3985 percent as massive as the carbon-12 atom.  Thus, if the mass of one carbon-12 atom is exactly 12 amu, the atomic mass of hydrogen must be 0.083985 × 12 amu, or 1.0078 amu. fluorine-19 is 18.9984 amu oxygen-16 is 15.9949 amu Similar calculations show that the atomic mass of fluorine-19 is 18.9984 amu and that of oxygen-16 is 15.9949 amu

59. 58 The average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes. The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance. Cont…

60. 59 Cont…

61. 60 2.2. Chemical Formulas  Chemists use chemical formulas to express the composition of molecules and ionic compounds in terms of chemical symbols.  By composition we mean not only the elements present but also the ratios in which the atoms are combined.  Here we are mainly concerned with two types of formulas: 1.molecular formulas 2.empirical formulas.

62. Cont… 61 A molecular formula:  shows the exact number of atoms of each element in the smallest unit of a substance.  From the molecules point of view, each example was given with its molecular formula in parentheses.  Thus, H 2 is the molecular formula for hydrogen, O 2 is oxygen, O 3 is ozone, and H 2 O is water.  The subscript numeral indicates the number of atoms of an element present in it.  There is no subscript for O in H 2 O because there is only one atom of oxygen in a molecule of water, and so the number “one” is omitted from the formula.  Note that oxygen (O 2 ) and ozone (O 3 ) are allotropes of oxygen. An allotrope is one of two or more distinct forms of an element.  Two allotropic forms of the element carbon—diamond and graphite—are dramatically different not only in properties but also in their relative cost.

63. 62 Empirical Formulas:  The empirical formula tells us which elements are present and are the simplest whole-number ratio of their atoms. Empirical formulas are the simplest chemical formulas; They are written by reducing the subscripts in the molecular formulas to the smallest possible whole numbers. Molecular formulas are the true formulas of molecules. If we know the molecular formula, we also know the empirical formula, but the reverse is not true. Cont…

64. 63  A molecule of metaldehyde (a pesticide used for snails and slugs) contains 8 carbon atoms, 16 hydrogen atoms, and 4 oxygen atoms. What are the molecular and empirical formulas of metaldehyde? Example 2.4 Empirical and Molecular Formulas Molecules of glucose (blood sugar) contain 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. What are the molecular and empirical formulas of glucose? Solution The molecular formula is C6H12O6 because one molecule actually contains 6C, 12H, and 6O atoms. The simplest whole-number ratio of C to H to O atoms in glucose is 1:2:1, so the empirical formula is CH2O  MF: C8H16O4  EF: 2:4:1 C2H4O Cont…

65. 64  For example, the molecular formula for acetic acid, the component that gives vinegar its sharp taste, is C2H4O2.  This formula indicates that a molecule of acetic acid contains two carbon atoms, four hydrogen atoms, and two oxygen atoms.  The ratio of atoms is 2:4:2. Dividing by the lowest common denominator (2) gives the simplest, whole-number ratio of atoms, 1:2:1, so the empirical formula is CH2O.  Note that a molecular formula is always a whole-number multiple of an empirical formula. Cont…

66. 65 2.3. The periodic table By the end of this section, you will be able to: State the periodic law and explain the organization of elements in the periodic table Predict the general properties of elements based on their location within the periodic table Identify metals, nonmetals, and metalloids by their properties and/or location on the periodic table

67. Periodic Table  Definition: Is a system for arranging the chemical element Mendeleev’s (1969) arrange elements in increasing of atomic mass However, In Modern periodic table elements arranged in of atomic number, usually in rows (Henry Moseley 1913) Has 7 periods and 18 groups 66 2.5. Historical development of the periodic table 66

68. 67 Before the beginning of the 18th century, it was easy to study and remember the properties of the elements because very few were known. However, in the middle of the 19th century, many more elements were discovered. Scientists then began to investigate possibilities for classifying them. After numerous attempts, the scientists were ultimately successful. They grouped elements with similar properties together. This arrangement is known as the classification of elements. Defects in Mendeleev's periodic table 1.Position of isotopes: Since elements are arranged in order of increasing atomic masses, the isotopes belong to different groups (because isotopes have different masses). 2. Wrong order of atomic masses of some elements: When certain elements are grouped on the basis of their chemical properties, some elements with higher atomic masses precede those with lower atomic masses. For example, argon, with atomic mass of 39.95, precedes potassium with atomic mass of 39.1. Cont…

69. 68 The Modern Periodic Law  Henry Mosley (In 1913): determined the atomic number of each of the elements by analyzing their X-ray spectra. 'No two elements can have the same atomic number.'  The atomic number of an element is the fundamental property that determines the chemical behavior of the element.  The discovery of atomic number led to the development of the modern periodic law.  The modern periodic law states that: ‘‘the properties of the elements are periodic function of their atomic numbers.’’  This means that when elements are arranged according to increasing atomic number, elements with similar physical and chemical properties fall in the same group. Cont…

70. 69 The periodic table organizes the elements in a particular way. A great deal of information about an element can be gathered from its position in the period table. For example, you can predict with reasonably good accuracy the physical and chemical properties of the element. You can also predict what other elements a particular element will react with chemically. Understanding the organization and plan of the periodic table will help you obtain basic information about each of the 118 known elements. Cont…  Many different forms of the periodic table have been published since Mendeleev's time. Today, the long form of the periodic table, which is called the modern periodic table, is commonly in use.  It is based on the modern periodic law.  In the modern periodic table, elements are arranged in periods and groups.

71. 70 Families Periods Columns of elements are called groups or families. Elements in each family have similar but not identical properties. For example, lithium (Li), sodium (Na), potassium (K), and other members of family IA are all soft, white, shiny metals. All elements in a family have the same number of valence electrons. Each horizontal row of elements is called a period. The elements in a period are not alike in properties. In fact, the properties change greatly across even given row. The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas. Cont…

72. Classification of the Elements in the Periodic Table  Classification of the elements in the periodic table can be done in three ways on the basis of their electronic configurations: 1. Representative elements: S-bock and P-block elementsP-block elements Elements in groups 1 and 2 are known as the s – block elements Group 13-18 are known as the p-block elements 2. Transition elements: Elements which belong to group 3 to 12 and have their outer shell electronic configuration as (n-1)d 1-10 ns 1-2. These elements are also known as the d-block elements. 3. Inner transition elements: Lanthanides and actinides series which fall at the bottom of the periodic tableperiodic table In these elements the 4f and 5f orbitals are partially filled, rendering them special properties. 71

73. 72  We can sort the elements into large classes with common properties: Metals: elements that are shiny, malleable, good conductors of heat and electricity; nonmetals and : elements that appear dull, poor conductors of heat and electricity; Metalloids: elements that conduct heat and electricity moderately well, and possess some properties of metals and some properties of nonmetals Cont…

74. 73 Cont… Periodic table diagram

75. 74 Periodic Properties within a Group 1.Nuclear Charge On moving down a given group of the periodic table, nuclear charge progressively increases, but effective nuclear charge remains nearly constant. 2. Atomic Size/Atomic Radius In moving down a group, atomic radius of the elements mainly depends on the number of shells. 3. Ionization Energy In which region of the periodic table do you find elements with the: a lowest tendency to lose electrons, and b highest tendency to lose electrons? Ionization energy is the minimum energy required to remove the outermost shell electron from an isolated gaseous atom or ion M (g) + energy –––– → M+ (g) + e– M (g) + energy –––– → M+ (g) + e– ⇒ First ionization energy M+ (g) + energy –––– → M2+ (g) + e– ⇒ Second ionization energy For a given element, the second ionization energy is higher than the first one. Generally, with the increasing atomic number, the first ionization energy decreases down the same group.

76. 75 4. Electronegativity Where do you find the most electronegative element in the periodic table? Electronegativity is the ability of an atom in a molecule to attract the shared electrons in the chemical bond. Fluorine, the most electronegative element and the least electronegative element is cesium. Cont…

77. 76 6. Metallic Character In which region of the periodic table do you find metals and non-metals? Metals have the tendency to lose electrons and form positive ions. As a result, metals are called electropositive elements. In moving down a group, atomic size increases progressively, and it becomes easier for elements to lose their valence electrons and form positive ions. Therefore, metallic character increases down a group. In the periodic table, metals and non-metals are separated by a stair step diagonal line, and elements near this border line are called metalloids. Metals are found on the left side of the line and nonmetals on its right side. Cont…

78. 77 Periodic Properties within a Period 1.Atomic Size: From left to right in a given period, nuclear charge or atomic number progressively increases by one for every succeeding element. However, increasing number of valence electrons is being added to the same shell. This results in an increase in effective nuclear charge. 2. Ionization Energy: Two factors account for the general trend in ionization energy across a given period. These are nuclear charge and atomic size. 3. Electron Affinity: The variation in electron affinity of elements in the same period is due to changes in nuclear charge and atomic size of the elements. A cross a period from left to right, electron affinity increases due to an increase in effective nuclear charge. The elements show a greater attraction for an extra added electron.

79. 78 4. Electronegativity: Across a period, a gradual change in nuclear charge and atomic size determine the trends in the electronegativity of the elements. 5. Metallic character: From left to right in a period, metallic character of the elements decreases. Elements on the left end of a period have a higher tendency to form positive ions. Those at the right end have a greater tendency to form negative ions. In any period, elements on the left side are metals and those on the right side are nonmetals. Cont…

80. 79 2.6. Ionic and Molecular Compounds  Compounds are classified as ionic or molecular (covalent) on the basis of the bonds present in them.  In ordinary chemical reactions, the nucleus of each atom (and thus the identity of the element) remains unchanged.  Electrons, however, can be added to atoms by transfer from other atoms, lost by transfer to other atoms, or shared with other atoms.  The transfer and sharing of electrons among atoms govern the chemistry of the elements.

81. 80  During the formation of some compounds, atoms gain or lose electrons, and form electrically charged particles called ions Metals loose electron---- Na -> Na+ + 1e- Non-metals gain electron Cl + 1e- -> Cl -  Ions can be: monatomic ion made from single atom eg:- Na+, Ca2+, Cl- etc.  Or polyatomic ions made from many atoms. eg:- SO 4 2-, NO 3 -, CO 3 2- etc. Cont…

82. 81 2.6.1.Ionic Compounds Formation  When an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of atoms that readily gain electrons (a non-metal), a transfer of electrons usually occurs, producing ions.  The compound formed by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions of opposite charge. eg:- formation of NaCl  A compound that contains ions and is held together by ionic bonds is called an ionic compound

83. 82 2.6.2.Molecular compounds Formation These molecular compounds (covalent compounds) result when atoms share rather than transfer (gains or losses), electrons; therefore, covalent bond is responsible for the formation of molecular compounds. H 2, Cl 2, CH 4  Properties of molecular compound- Under normal conditions, molecular compounds often exist as gases, Low-boiling liquids, Low-melting solids covalent compounds are usually formed by a combination of nonmetals.

84. 83 Eg. Predict whether the following compounds are ionic or molecular(covalent)? (a) KI, the compound used as a source of iodine in table salt (b) H 2 O 2, the bleach and disinfectant hydrogen peroxide (c) CHCl 3, the anesthetic chloroform (d) Li 2 CO 3, a source of lithium in anti depressants Cont…

85. 84 2.7.Chemical nomenclature  Nomenclature, a collection of rules for naming things, is important in science 2.7.1. Ionic compounds 2.7.1.1. Compounds Containing Only Monatomic Ions  The name of a binary compound containing monatomic ions consists of the name of the cation followed by the name of the anion (the name of the nonmetallic element with its ending replaced by the suffix -ide).

86. 85 2.7.1.2. Compounds Containing Polyatomic Ions  Compounds containing polyatomic ions are named similarly to those containing only monatomic ions, i.e. by naming first the cation and then the anion. Cont…

87. 86 2.7.1.3. Compounds Containing a Metal Ion with a Variable Charge Most of the transition metals can form two or more cations with different charges. Compounds of these metals with nonmetals are named as the charge of the metal ion is specified by a Roman numeral in parentheses after the name of the metal Cont…

88. 87 2.7.1.4. Ionic Hydrates  Ionic compounds that contain water molecules as integral components of their crystals are called hydrates.  The name for an ionic hydrate is derived by adding a term to the name for the anhydrous compound with each formula unit of the compound. For example, the anhydrous compound copper(II) sulfate also exists as a hydrate containing five water molecules and named copper(II) sulfate pentahydrate (CuSO 4.5H 2 O). Cont…

89. 88 Cont…

90. 89 2.7.2. Molecular (Covalent) Compounds  The bonding characteristics of inorganic molecular compounds are different from ionic compounds and they are named using a different system as well. 2.7.2.1. Compounds Composed of Two Elements  When two nonmetallic elements form a molecular compound, several combination ratios are often possible. For example, carbon and oxygen can form the compounds CO and CO 2. Cont…

91. 90 The numbers of atoms of each element are designated by the Greek prefixes. When only one atom of the first element is present, the prefix mono- is usually deleted from that part. Thus, CO is named carbon monoxide, and CO2 is called carbon dioxide. When two vowels are adjacent, the a in the Greek prefix is usually dropped. Cont…

92. 91 Cont…

93. 92  If the compound is a binary acid (comprised of hydrogen and one other nonmetallic element): 1. The word “hydrogen” is changed to the prefix hydro- 2. The other nonmetallic element name is modified by adding the suffix -ic 3. The word “acid” is added as a second word For example, when the gas HCl (hydrogen chloride) is dissolved in water, the solution is called hydrochloric acid. Names of Some Simple Acids 2.7.2.2. Binary Acids Cont…

94. 93 C ON …  Oxyacids are compounds that contain hydrogen, oxygen, and at least one other element, and are bonded in such a way as to impart acidic properties to the compound.  Typical oxyacids consist of hydrogen combined with a polyatomic, oxygen-containing ion. To name oxyacids: 1. Omit “hydrogen” 2. Start with the root name of the anion 3. Replace –ate with –ic, or –ite with –ous 4. Add “acid” 93 2.7.2.3. Oxyacids

95. CHAPTER 3 COMPOSITION OF SUBSTANCES AND SOLUTIONS 94

96. 3.1.1. Formula Mass Formula Mass for Covalent Substances 3.1. Formula Mass and Mole Concept Formula mass of a substance calculated by summing the average atomic masses of all the atoms represented in the substance’s formula For covalent substances, the formula represents the numbers and types of atoms composing a single molecule of the substance; therefore, formula mass may be correctly referred to as a molecular mass. 95

97. Cont…  Consider chloroform (CHCl 3 ), a covalent compound once used as a surgical anesthetic and now primarily used in the production of tetrafluoroethylene, the building block for the "anti-stick" polymer, Teflon.  What is the molecular mass (amu) for chloroform (CHCl 3 ) 96

98. Cont… Formula Mass for Ionic Compound Consider sodium chloride, NaCl, its molecular mass is 58.44 Computing Formula Mass for an Ionic Compound, Aluminum sulfate, Al 2 (SO 4 ) 3, formula mass is 342.14 amu 97

99. 3.1.2.The Mole Concept atoms or ions  The identity of a substance is defined not only by the types of atoms or ions it contains, but by the quantity of each type of atom or ion.  The mole is an amount unit, similar to familiar units like pair, dozen, gross, etc.  It provides a specific measure of the number of atoms or molecules in a sample of matter.  A mole of substance contains 6.02214076 X 10 23 discrete entities (atoms or molecules). 98

100. 99  The molar mass of an element (or compound) is the mass in grams of 1 mole of that substance Avogadro’s number (NA)  This large number is a fundamental constant known as Avogadro’s number (NA) or the Avogadro constant (Amedeo Avogadro).  This constant is properly reported with an explicit unit of “per mole,” a conveniently rounded to 6.022 X 10 23 /mol. Cont…

101. Practice Exercise 1.According to nutritional guidelines from the US Department of Agriculture, the estimated average requirement for dietary potassium is 4.7 g. What is the estimated average requirement of potassium in moles? Ans. 0.12 mol K 2.A liter of air contains 9.2 × 10 −4 mol argon. What is the mass of Ar in a liter of air? Ans. 0.0367g Ar 3. Copper is commonly used to fabricate electrical wire. How many copper atoms are in 5.00 g of copper wire? Ans.0.0787 mol of Cu 4. Beryllium is a light metal used to fabricate transparent X-ray windows for medical imaging instruments. How many moles of Be are in a thin-foil window weighing 3.24 g? Ans. 0.36 mol of Be 100

102. 101 Example. Deriving the Number of Atoms and Molecules from the Mass of a Compound A packet of an artificial sweetener contains 40.0 mg of saccharin (C 7 H 5 NO 3 S), which has the structural formula: Given that saccharin has a molar mass of 183.18 g/mol, how many saccharin molecules are in a 40.0-mg (0.0400-g) sample of saccharin? How many carbon atoms are in the same sample? Solution : first find mole, then multiply it with Avogadro number 101 Cont…

103. 3.2. Determining Empirical and Molecular Formulas  From chemical formula of the substance, one may determine the amount of the substance (moles) from its mass, and vice versa  But what if the chemical formula of a substance is unknown? In this section, these same principles will be applied to derive the chemical formulas of unknown substances from experimental mass measurements. 3.2.1. Percent Composition Measuring the mass of each of its constituent elements permit the calculation of the compound’s percent composition, defined as the percentage by mass of each element in the compound. 102

104. Cont… Example 1: Analysis of a 12.04-g sample of a liquid compound composed of carbon, hydrogen, and nitrogen showed it to contain 7.34 g C, 1.85 g H, and 2.85 g N. What is the percent composition of this compound?  To calculate % composition, divide the experimentally derived mass of each element by the overall mass of the compound, and then convert to a percentage: %=1.85/(12.04 ) 100%=15.4% %=2.85/(12.04 ) 100%=23.7% %=7.34/(12.04 ) 100%=61.0% Solution 103

105. Cont… 104

106. Cont… Determining Percent Composition from Molecular or Empirical Formulas Example 2  Determining Percent Composition from a Molecular Formula Aspirin is a compound with the molecular formula C 9 H 8 O 4. What is its percent composition of each atom? Solution: %C =(9x12/244.14)x100 % = (108.09/244.14)x100= 44.275 % %H =(8x1.008/244.14)x100 % = (8.046/244.14)x100= 3.295 % %O =(4x32/244.14)x100 % = (128/244.14)x100= 52.428 % 105

107. Cont… 3.2.2. Determination of Empirical Formulas (a)From mass of element Tools/steps used for determining empirical formula of a given compound 106

108. 107 What is the empirical formula of a compound if a sample contains 0.130 g of nitrogen and 0.370 g of oxygen? Determining an Empirical Formula from Percent Composition per the definition for percent composition, the mass of a given element in grams is numerically equivalent to the element’s mass percentage.

109. Cont… Example: Cisplatin, the common name for a platinum compound that is used to treat cancerous tumors, has the composition (mass percent) 65.02% platinum, 9.34% nitrogen, 2.02% hydrogen, and 23.63% chlorine. Calculate the empirical formula for cisplatin? Solution Where do we want to go? Empirical formula _ P t a N b H c C l d a----? b-----? c-------? d--------? What do we know? The sample of compound contains: 65.02% Pt 9.34% N 2.02% H 23.63% Cl Atomic masses Pt = 195.1 g/mol H = 1.008 g/mol N = 14.01 g/mol Cl = 35.45 g/mol 108

110. Cont… 109

111. Cont…  The empirical formula for cisplatin is Pt N 2 H 6 Cl 2. Note that the number for hydrogen is slightly greater than 6, because of rounding it become 6.. 110

112. 111 What is the empirical formula of a compound containing 40.0% C, 6.71% H, and 53.28% O? Answer: CH2O

113. Cont… 3.2.3. Determination of molecular formulas  Molecular formulas are derived by comparing the compound’s molecular or molar mass to its empirical formula mass.  If the molecular (or molar) mass of the substance is known, it may be divided by the empirical formula mass to yield the number of empirical formula units per molecule (n): So generally  For any empirical formula AxBy, the molecular formula is determined by multiplying it to whole number n, Where n is ratio 112

114. Cont… Guide to calculating a molecular formula from an empirical formula: Obtain the empirical formula and calculate the empirical formula mass. Divide the molar mass by the empirical formula mass to obtain a small integer. Multiply the empirical formula by the small integer to obtain the molecular formula. Example: Melamine, which is used to make plastic items such as dishes and toys, contains 28.57% C, 4.80% H, and 66.64% N. If the experimental molar mass is 125 g, what is the molecular formula of melamine? (Basic Chemistry). 113

115. Cont… Example-1 A compound whose empirical formula is determined to be CH 2 O. If the compound’s molecular mass is determined to be 180 amu by Mass spectrometry, what is the molecular formula of the compound? Solution First find molecular mass of the emperical formula CH 2 O, it is 30 (12C+2H+16’O’). Then find the whole number n n = molecular mass/emperical formula mass= 180/30=6 Then molecular formula is (CH 2 O)n= (CH 2 O) 6 = C 6 H 12 O 6 Example-2 Determination of the Molecular Formula for Nicotine Nicotine, an alkaloid in the nightshade family of plants that is mainly responsible for the addictive nature of cigarettes, contains 74.02% C, 8.710% H, and 17.27% N. If 40.57 g of nicotine contains 0.2500 mol nicotine, what is the molecular formula? 114

116. Cont… 115

117. Cont… 116

118. Cont… 117

119. 118 Whatisthemolecularformulaofacompoundwithapercentcompositionof49.47%C,5.201 %H,28.84% N, and 16.48% O, and a molecular mass of 194.2 amu? Answer: C8H10N4O2

120. 3.3. Molarity and Other Concentration Units 3.3.1. Molarity  Solutions are defined as homogeneous mixtures  The relative amount of a given solution component is known as its concentration.  Solvent may be viewed as the medium in which the other components are dispersed, or dissolved.  A solution in which water is the solvent is called an aqueous solution.  A solute is a component of a solution that is typically present at a much lower concentration than the solvent.  Solute concentrations are often described with qualitative terms such as:  dilute (relatively low concentration) and  Concentrated (relatively high concentration ). 119

121. Cont… Example: Distilled white vinegar is a solution of acetic acid, CH 3 CO 2 H, in water. A 0.500-L vinegar solution contains 25.2g of acetic acid. What is the concentration of the acetic acid solution in units of molarity? 120 What volume of a 1.50-MKBr solution contains 66.0 g KBr? Answer: 0.370 L

122. 3.3.2.Dilution of Solutions Dilution is the process where by the concentration of a solution is lessened by the addition of solvent. Dilution is also a common means of preparing solutions of a desired concentration. By adding solvent to a measured portion of a more concentrated stock solution. Guide to Calculating Dilution Quantities 1. Prepare a table of the concentrations and volumes of the solutions. 2. Rearrange the dilution expression to solve for the unknown quantity. 3. Substitute the known quantities into the dilution expression and calculate. 121

123. Cont… Example: If 0.850 L of a 5.00-M solution of copper nitrate, Cu(NO 3 ) 2, is diluted to a volume of 1.80 L by the addition of water, what is the molarity of the diluted solution? Solution:- 122

124. 3.3.3. Percentage (W/W, W/V and V/V) 123

125. Cont… 124

126. Cont… Parts per million (ppm) and Part per billion (ppb)  Very low solute concentrations are often expressed using appropriately small units such as parts per million (ppm) or parts per billion (ppb).  The mass-based definitions of ppm and ppb are given here: 125

127. The End of CH-3 Thanks! 126