Atomic Structure & Basic Concepts of Chemistry PDF

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See discussions, Atomic Structure & Basic Concepts Chapter · April 2018 CITATIONS 0 1 author: Ramesh Duraisamy Arba Minch University 34 PUBLICATIONS 262 CITATIONS SEE PROFILE Some of the authors of this publication A Review of Bio-tanning Materials Fish Leather processing and All content following this page was uploaded The user has requested enhancement of E – MATERIAL for CHEMISTRY GRADUATES Atomic Structure & Basic Concepts Prepared By Dr. RAMESH DURAISAMY Assistant Professor Department of Chemistry Arba Minch University Arba Minch (Ethiopia) Chapter. I Introduction to Chemistry..............4 1.1 Suggestions for Studying General Chemistry, practicing with the new fact or skill....................4 1.1.1 Take responsibility for your own learning 1.1.2 Studying............................. 1.1.3 Develop an effective set of study skills...4 1.1.4 Develop your problem-solving skills......4 1.2 The Scientific Methods................... 1.2.1 Law and Hypothesis.................. 1.2.2 Theory.............................. 1.3 Measurement and units in chemistry.............5 1.3.1 SI units............................ 1.3.1.2 Density and Specific gravity................6 1.3.1.4 Mole concept........................ 1.4 The Classification of Matter............ 1.4.1 Three states of matter.............. 1.5 Some Basic Definitions.................. 1.5.1 Physical and Chemical properties...........8 1.5.2 Atoms and Molecules.............................8 1.5.3 Substances and Mixtures........................9 1.5.4 Elements and Compounds......................9 1.5.5 Molecules and Ions............................... 1.5.5.1 Molecules........................... 1.6 Significant Figures..................... 1.6.1 Guidelines for using Significant figures 10 Chapter-II Atomic Structure........ 2.1 Atomic theory & fundamental particles....... 12 Atomic Theory................................ 2.2 The structure of the Atom........................... 2.2.1 The Electron........................ 2.2.2 The proton and Nucleus........................ 2.2.3 The Neutron......................... 2.3 Atomic number, Mass number and Isotopes 14 2.4 Nature of light and Electromagnetic radiation 2.4.1 Corpuscular Theory.............................. Chapter-III Periodic 3.1 Need to classify elements............................ 3.2 Brief history of modern periodic law and modern periodic table................... 3.2.1 Modern periodic Law........................... 3.2.2 Modern periodic table ( Long form of... 28 periodic table).......................... 3.3 Class of periodic elements : Based on electronic configuration.................... 3.3.1 s - block elements................. 3.4 Chemical Properties of s-block 3.4.1 IA Group elements (Alkali 3.5 p-block elements........................ 3.5.1 Introduction....................... 3.5.2 General aspects and trends elements................................. 3.6 Groups in p-block elements......................... 3.6.1 Boron Family....................... 3.6.2 Elements of group IV A (Carbon 3.6.3 Elements of VA (15th) group family).................................. 3.6.4 Oxygen family (Elements of 3.7 IIA Group elements (Alkaline earth Chapter - IV Molecular 4.1 Molecular Geometry................................... 4.2 Dipole moments.......................... 4.3 Valence Bond Theory................................. Chapter. V Chemical Reactions.................. 5.1 Chemical Equations.................................... 5.1.1 Writing chemical equations.................. 5.2 The mole concept........................ 5.3 Balancing of the chemical equations........... 5.4 Amount of reactants and products............... 5.5 Limiting Reagents....................... 5.6 Reaction Yield.......................... Atomic Structure & Basics of Chemistry Chapter. I Introduction to Chemistry 1.1 Suggestions for Studying General Chemistry One of the most common challenges that General Chemistry students have is that they don't realize the importance of being actively involved in their own learning. Many seem to think that coming to class and passively listening to the instructor's pearls of wisdom is all that is needed. Generally, a little different; everyone has at least different intellectual equipment and a life and school experiences, which they learning situation. Therefore, the learning experience is a little different for each person, even those in same classroom. The person who wishes to become a learner must be actively involved in constructing his or her own set of ideas. This process will be different in its details for each person, but each must go through a similar of hearing (or seeing or reading) the new item, comparing it with what is already in his and integrating the new information into the memory bank. This is not a simple process--it requires work! The work may take the form of reading and re-reading, thinking about the new information and how it fits in, practicing with the new fact or skill. 1.1.1 Take responsibility for your own learning  Put in the required effort  Read the assignments carefully before coming to class, and then again after class.  Do the assigned problems and may be extras for good measures.  Come to class prepared and on time. Class attendance is important even if your instructor doesn't specifically require it.  Take complete notes in class  Pay attention in class. Listen to your instructor, not to the student next to you.  If you miss class, make sure that you get all the necessary information from the instructor or another responsible student.  Try to capture the ideas & concepts Atomic Structure & Basics of Chemistry 1.2 The Scientific Methods The study of chemistry concerned with the observations, theories and laws that give this science its foundation and that form a framework into which information fits to make an integrated area of knowledge. This framework develops from the persuit of answers to many questions, can be subjected to experimental investigation by an approach often called the scientific method. 1.2.1 Law and Hypothesis Scientists identify problems or questions through their own observations and experiments and get the conclusions, compare with earlier observations of the same. The first step of the scientific method, to a problem involves carefully planning the experiments to gather the information about all phases of the problem. The results are examined for general relationships that will unify the observations. Sometimes a wide variety of observations can be summarized in a general verbal statement or mathematical equation known as a law. One example is the law of conservation of mass, which summarizes the results of thousands of experimental observations. More often, is suggested a tentative explanation. Such a proposal hypothesis. 1.2.2 Theory A hypothesis is tested by further experiments, and, if it is capable of explaining the large body of experimental data, it is dignified by the name of theory. The data obtained in a research study may be both quantitative, consisting of general observations about the system, and qualitative, comprising numbers obtained by various measurements and observations. 1.3 Measurement and units in chemistry The measurements chemists make are often calculations to obtain other related quantities Different instruments enable us to measure a substance’s properties. The meter stick measures length or scale; the buret, the pipette, the graduated cylinder and volumetric flask measure volume; the balance measures mass; the thermometer measures temperature. These instruments measurements of microscopic properties, which can Atomic Structure & Basics of Chemistry Chemists are interested primarily in mass, which can be determined readily with a balance; the process of measuring mass is called weighing. The SI base unit of mass is the kilogram (kg), but in chemistry the smaller gram (g) is more convenient; 1 kg = 1000g = 1x103g 1.3.1.1 Mass and Weight Mass is a measure of the amount of matter in an object, it is an invariable quantity. On the other hand, the weight of a body is the force that gravity exerts on the body; it is variable, since the attraction depends on the distance from planet’s center. Scientists measure quantities of matter in terms of mass rather than weight because the mass remains constant, whereas its weight is of its environment. However, that the often used loosely for the mass of a substance. The mass of an object is found by comparing it to other objects of known mass. The instrument used to make this comparison is called a balance. The standard object to which all SI and other metric units of mass are referred a cylinder of platinum-iri alloy. The unit of mass, the kilogram (kg), is defined as the mass of this cylinder; some times kilogram expressed in pounds (1 kilogram is about 2.2 pounds). The gram (g) is equal to exactly 0.001 times mass of the std kilogram (1x10-3 kg) and is nearly equal to the weight of 1cm3 of water. 1.3.1.2 Density and Specific gravity One of the physical properties of a solid, liquid or a gas is its density. Density is defined as mass per unit volume. This is mathematically expressed as, Mass M Density = ------------- (or) Volume V Substances usually have different densities, Substances Density (at 250C and 1atm.) Air 1.29 g/L He 0.179 g/L H2O 0.997 g/cm3 Glycerin 1.26 g/cm3 Hg 13.6 g/cm3 Salt 2.17 g/cm3 Iron 7.86 g/cm3 Silver 10.5 g/cm3 Here, density for all cases expressed as, g/cm3 Atomic Structure & Basics of Chemistry Workout.3 A piece of gold in got with a mass of 301g has a volume of 15.6 cm3. Calculate the density of gold. D = M/V = 301g / 15.6 cm3 D = 19.3 g/cm3 1.3.1.3 Volume The SI unit of length is the meter (m) derived unit for volume is the cubic meter (m3). Generally, chemists work with smaller volumes, such as cubic centimeter (cm3) and the cubic decimeter (dm3). 1cm3 = (1x10-2m)3 = 1x10-6m3 1dm3 = (1x10-1m)3 = 1x10-3m3 Another common unit of volume is the liter (L). A liter is the volume occupied by one cubic decimeter. One liter volume is equal to 1000 mL or 1000 cm3. 1L = 1000 mL = 1000 cm3 (1mL = 1cm3) = 1 dm3 1.3.1.4 Mole concept Usually, number of atoms is measured in the units of 6.023x1023. This unit called one mole of atoms (1mol). The number of atoms in one mole of any element (1 mol. of atom has 6.023 x 1023 atoms) is also called Avogadro’s number. For example, let us consider a sample of carbon tetra chloride by combining 1 mole of carbon atoms with 4 moles of Cl atoms. This would be a produce 6.023 x 1023 molecule of CCl4. Workout.4 How many CCl4 molecules can be made by the combination of 0.050 mol of carbon atoms with Cl atom? The formula of CCl4 tells us that one CCl4 molecule contains one carbon atom (1 molecule CCl4/ 1 atom of C). If we can determine number of carbon atoms present, we can determine the number of CCl4 molecules that can be made. One mole of C atoms, contains 6.023x1023 C atoms; 0.050 mole of C atoms contains, 0.050 mol C atoms x 6.023 x 1023 mol of C atoms ----------------------------------------------- 1 mol of C atoms = 3x1022 C atoms A single CCl4 molecule contains 1 atom Therefore, we can make 1 molecule of CCl4 Atomic Structure & Basics of Chemistry together in an orderly fashion with little freedom of motion. Molecules in a liquid are close together but are not held so rigidly in position and can move past one another. In a gas, distances that are large compared with the size of the molecules separate the molecules. The three states of matter can be interconverted without changing the composition of the substance. Upon heating, a solid (ice) will melt to form a liquid (water) the temperature at which this transition occurs is called the melting point. Further heating will convert the liquid into a gas, this conversion take place at the boiling point of the liquid. On the other hand, cooling a gas will cause it to condense into a liquid. When the liquid is cooled further, it will freeze into the solid form. 1.5 Some Basic Definitions: Mass and Weight, Physical and Chemical properties, Atoms and molecules, Substances and Mixtures, Elements and Compounds. 1.5.1 Physical and Chemical properties Substances are identified by their properties as well as composition. Color, melting point, and boiling point are physical properties. A physical property can be measured and observed without changing the composition or identity of a substance. example, we can measure the melting point heating a block of ice and recording the at which the ice is converted to water. from ice only in appearance, not in composition, this is a physical change; we can freeze recover the original ice. Therefore, the of a substance is a physical property. when we say that helium gas is lighter are referring to a physical property. On the other hand the statement. “Hydrogen burns in oxygen gas to form water”, chemical property of hydrogen, because observe this property we must carry out change, in this case burning. After the original chemical substance, the hydrogen have vanished, and all that will be left chemical substance. In the case, of water recover the hydrogen from the water by physical change, such as boiling or freezing. Atomic Structure & Basics of Chemistry quite different from those of water and those of each other. 1.5.3 Substances and Mixtures A substance is a form of matter that has a definite (constant) composition and distinct properties. Examples are water, ammonia, sucrose, gold and oxygen. Substances differ from one another in composition and can be identified by their appearance, smell, taste, and other properties. A mixture is a combination of two or more substances in which the substances retain their distinct identities. Some familiar examples are air, soft drinks, milk, and cement. Mixtures do not have constant composition. Mixtures are either homogeneous or heterogeneous. When a spoonful of sugar dissolves in water we obtain a homogeneous mixture in which the composition of the mixture is the same throughout. If sand is mixed with iron filings, however, the sand grains and the iron filings remain separate. This type of mixture is called a heterogeneous mixture because the composition is not uniform. Any mixture, whether homogeneous heterogeneous, can be created and then separated by physical means into pure components without changing the identities of the components. sugar can be recovered from a water solution by heating the solution and evaporating it Condensing the vapor will give us back component. 1.5.4 Elements and Compounds An element is a substance that cannot be separated into simpler substances by chemical means. To date, 115 elements have been positively identified. Eighty-three of them occur naturally on Earth. Scientists via nuclear processes have created the others. Only a few elements, such as the gases helium, neon, and argon, consist of a collection of individual atoms that move about independently of one another. Other elements, such as the gases nitrogen, oxygen and chlorine, consist of pairs of atoms, each pair moving as a single unit. The element phosphorous consists of units composed of four phosphorus atoms; sulfur, of units composed of eight sulfur atoms. Atomic Structure & Basics of Chemistry 1.5.5 Molecules and Ions Out of all the elements, only the six noble gases (He, Ne, Ar, Kr,Xe, and Rn) exist in nature as single atoms. For this reason, they are called monatomic (mean a single atom) gases. Most matter is composed of molecules or ions formed by atoms. 1.5.5.1 Molecules A molecule is formed from least two atoms in a definite arrangement held together by chemical forces (called chemical bonds). A molecule may contain atoms of the same element or atoms of two or more different elements joined in a fixed ratio. Thus, a molecule is not necessarily a compound, which is made up of two or more elements. For example hydrogen gas is a pure element, but it consists of molecules made up of by two H atoms. On the other hand, water is a molecular that contains hydrogen and oxygen in a ratio of two H atoms and one O atom. Like atoms, molecules are electrically neutral. The hydrogen molecule, H2, is called a diatomic molecule because it contains only two atoms of the same element. Other elements that normally exist as diatomic molecule are nitrogen (N2) and oxygen (O2) as well as F2, Cl2 and I2. A diatomic molecule can also contain atoms of different elements. For example are hydrogen chloride (HCl) and CO. The vast majority of molecules contain more than two atoms. They can be atoms of the same element, as in ozone (O3), which is made up of three atoms of oxygen, or they can be combinations of two or more different elements. Molecules containing more than two atoms are called polyatomic molecules. Like ozone, water (H2O) and ammonia (NH3) are polyatomic molecules. 1.5.5.2 Ions An ion is an atom or a group of atoms that has a net positive or negative charge. The loss of one or more electrons from a neutral atom results a cation (positive charge). For example, a sodium atom (Na) can readily lose an electron to become sodium cation, which is represented by Na+: Na atom: 11 protons and 11 electrons Na+ ion : 11 protons and 10 electrons An anion is an ion whose net charge is negative due to an increase in the number of electrons. Atomic Structure & Basics of Chemistry  If a number is greater than 1, then all the zeros written to the right of the decimal point count as significant figures. Thus 2.0mg has two significant figures, 40.062 mL has five significant figures, and 3.040 dm has four significant figures. If a number is less than 1, then only the zeros that are at the end of the number and the zeros that are between nonzero digits are are significant. This means that 0.090 kg has two significant figures, 0.3005 L has four significant figures, 0.00420 min. has three significant figures, and so on.  For numbers that do not contain decimal points, the trailing zeros may or may not be significant. Thus 400 cm may have one significant figure (the digit 4), two significant figures (40), or three significant figures (400). We cannot know which is correct without ambiguity. In this particular case, we can express the number 400 as 4x102 for one significant figure, 4.0x102 for figures, or 4.00x102 for three significant figures. The following example shows the determination of significant figures. Example.1 Determine the number of significant figures in the following measurements: (a) 468 cm (b) 5.03 g (c) 0.798 m (d) 0.046 kg (e) 23 1.340x10 atoms (f) 6000ml Solution: We follow the rules for determining significant figures. (a) Three (b) Three (c) Three (d) Two (e) Four (f) This is an ambiguous case. The number of significant figures may be four (7.000x103), three (7.00x103), two (7.0x103), or one (7x103). A second set of rules specifies how to handle significant figures in calculations.  In addition and subtraction, the answer have more digits to the right of the than either of the original numbers. these examples. 89.332 + 1.1 = 90.432 -------- round 2.097 – 0.12 = 1.977 --------- round The rounding off procedure is as follows, to round off a number at a certain point drop the digits that follow if the less than 5. Thus, 8.723 rounds of Atomic Structure & Basics of Chemistry Chapter-II Atomic Structure 2.1 Atomic theory & fundamental particles Atomic Theory In fifth century, the Greek philosopher Democritus expressed that all matter consists of very small, indivisible particles, named as ‘atomos’ (means uncuttable or indivisible), but his idea was not accepted. Experimental evidence from early scientific investigations provides support for the ‘atomism’ and gradually gave the modern definitions of elements and compounds. It was in 1808, English scientist John Dalton, formulated a precise definition of the indivisible building blocks of matter that we call atoms. According to this Daltons’s summarized the atomic theory as, 1. Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. 2. Compounds are composed of atoms of more than one element. In any compound, the ratio of the number of atoms of any two elements either an integer or a simple fraction. 3. A chemical reaction involves only the separation, combination or rearrangement. Dalton’s concept of an atom was far more detailed and specific than Democritus and Dalton’s proposed the three hypotheses are as follows, First hypothesis It states that atoms of one element are different from atoms of all other elements. Dalton made no attempt to describe the structure or composition of atoms, he had no idea about, what an atom is really like. But he did realize that the different properties shown by different elements such as hydrogen and oxygen can be explained by assuming that hydrogen atoms are not the same as oxygen atoms. Second hypothesis The 2nd hypothesis suggests that, in order to form a certain compound, we need not only atoms of Atomic Structure & Basics of Chemistry passes through a hole and continues traveling to the other end of the tube. When the ray strikes the specially coated surface, it produces a strong fluorescence, or bright light. In some experiments two electrically charged plates and a magnet were added to the outside of the cathode ray tube. When the magnetic field is on and the electric field is off, the cathode ray strikes point A. When only the electric field is on, the ray strikes point C. When both magnetic and the electric fields are off or when they are both on but balanced so that they cancel each other’s influence, the ray strikes point B. According to electromagnetic theory, a moving charged body behaves like a magnet and can interact with electric and magnetic fields through which it passes. Since the cathode ray is attracted by the plate bearing positive charges and repelled by the plate bearing negative charges, it must consist of negatively charged particles. We know these negatively charged particles as electrons. An English physics, J.J.Thomson, used a cathode ray tube and his knowledge of electromagnetic theory to determine the ratio of electric charge to the mass of an individual electron. The number he came up, -1.76x108 C/g, where C stands for coulomb, which is the unit of electric charge. Thereafter, in a series of experiments carried out, R.A.Millikan succeeded in measuring the charge of the electron with great precision. His work proved that the charge on each electron was exactly the same. In his experiment, he suspended the charged drops in air by applying an electric field and followed their motions through a microscope. Using his knowledge of electrostatics, Millikan found the charge of an electron to be –1.6022x10-19 C. From these data he calculated the mass of an electron. Charge mass of an electron = ------------------- Charge/mass -1.6022x10-19C = --------------------- = 9.10 x 10-28 g - 1.76 x 108 C/g Radioactivity The German physicist Wilhelm Rontgen noticed 1895), that cathode rays caused glass Atomic Structure & Basics of Chemistry experiments using very thin foils of gold and other metals as targets for  particles from a radioactive source. He was observed that the majority of particles penetrated the foil either un-deflected or with only a slight deflection. But, a  particle was scattered at a large angle. Rutherford, explained the results of the -scattering experiment in terms of a new model for the atom. According to Rutherford concept, most of the atom must be an empty space. This explains why the majority of  particles passed through the gold foil with little or no deflection. The atoms have positive charges, which is in central core of the atom, called nucleus. The positively charged particles in the nucleus are called protons. In separate experiments, it was found that each proton carries the same quantity of charge as an electron and has a mass of 1.6262 x 10-24 g about 1840 times the mass of the oppositely charged electron. 2.2.3 The Neutron Rutherford’s model of atomic structure left one major problem unsolved. It was known that hydrogen, the simplest atom, contains only one proton and that the helium atom contains two protons. Therefore, the ratio of the mass of a helium atom to that of a hydrogen atom should be 2:1. In reality, however, the ratio is 4:1. Rutherford and others postulated that there must be another type of subatomic particle in the atomic nucleus; another English physicist, James Chadwick, provided the proof. When Chadwick bombarded a thin sheet of beryllium with  particles, a very high-energetic radiation similar to  rays was emitted by the metal. Later experiments showed that the rays actually consisted of a third type of subatomic particles, which Chadwick named neutrons, because they proved to be electrically neutral particles having a mass (1.67493 x 10-24), slightly greater than that of protons. 2.3 Atomic number, Mass number and Isotopes N Atomic number The number of protons and neutrons they contain can identify all atoms. The atomic number number of protons in the nucleus of each element. In a neutral atom the number Atomic Structure & Basics of Chemistry some facts such as reflection and refraction. But it is failed to explain interference and diffraction. Hygens proposed wave like character of light explains the phenomenon of interference and diffraction. In 1856 James Clark Maxwell proposed that light in the form of waves. These waves have electric and magnetic field associated with them, therefore electromagnetic radiations or electromagnetic waves. 2.4.2 Characteristic of electromagnetic radiation 1. These consist of electric and magnetic fields that oscillate in the directions perpendicular to each other. 2. All electromagnetic waves travel with same velocity. The velocity equal to that of light, 3x108 m/sec. 3. These electromagnetic radiations donot require any medium for propagation. E.g. light us from the Sun through an empty space 2.4.3 Properties of Waves and Electromagnetic radiation According to Rutherford’s model, an atom consists of a nucleus is many times smaller than the atom itself, with electrons occupying the remaining space. Each element has a characteristic line spectrum because of the emission of light from atom in the hot gas. The spectra can be used to identify elements. This spectrum of each element was characterized through the waves (i.e. electromagnetic radiation). Wave motion A familiar example of a electromagnetic waves like waves in the surface of water. The waves originate from the centre of disturbance and propagate in the form of up and down movements. That is, if we drop a stone into one end pond, the impact of the stone with water starts an up-and-down motion of the water surface. This up- and-down motion travels outward from where the stone hit. A wave is a continuously repeating change or oscillation in matter at regular intervals. Light is also a wave; it consists of oscillations in electric and magnetic fields that can travel through space. Visible light, X rays and radio waves all are the Atomic Structure & Basics of Chemistry 2.5 Electromagnetic Radiation There are many kinds of waves, such as water waves, sound waves, and light waves. In 1873 James Clerk Maxwell proposed that light consists of electromagnetic waves. According to Maxwell’s theory, an electromagnetic wave has an electric field component and a magnetic field component. These two components have the same wavelength and frequency, and hence the same speed, but they travel in mutually perpendicula 10-12 10-11 10-10 10-8 10-7 Wavelength (nm)  rays X rays UV UV Frequency (Hz) 1020 1018 1016 Fig.1a Violet 400 nm 500nm Fig.1b Visible light The fig.1a shows various types of electromagnetic radiation, which differ from one another in wavelength and frequency. The range of frequencies or wavelengths of electromagnetic radiation is called the electromagnetic spectrum. The motions of electrons within atoms and molecules produce the shorter, visible light waves. The shortest waves, which also have the highest frequency, are associated with  rays, which result from changes within the nucleus of the atom. In fig.1b visible light extends from the violet end of the spectrum, which has a wavelength of about 400nm, to the red end, with a wavelength of less than 800 nm. Beyond these extremes, electromagnetic radiation is not visible to the human eye. IR radiation has wavelengths greater Atomic Structure & Basics of Chemistry 2.6.1 Planck’s Quantization Energy Max Planck found a theoretical formula that exactly describes the intensity of light of various frequencies emitted by a hot solid at different temperatures. According to Planck, the atoms of the solid oscillate, or vibrate with a definite frequency,, depending on the solid and he found it necessary strong idea. 1. Radiant energy is not emitted or absorbed continuously but discontinuously in the form of small pockets of energy called quanta. Each such quanta is associated with a definite amount of energy. In the case of light, the quanta of energy are often called photon. 2. The amount of energy associated with a quantum of radiation is proportional to the frequency of light. E or E = h and Where, h is a Planck’s constant with the value 6.63 x 10-34 J.s. This relation found to be valid for all type of electromagnetic radiation. 3. The total amount of energy emitted or absorbed by a body (matter) will be some whole number multiple of quantum. E= nh ----------- (1) n=1,2,3,………… This means that a body can emit or absorb energy equal to h, 2h, 3h, ………. Or any other integral multiple of h. But cannot emit or absorb fractional value of h. We know that, C = ; = C/ E= h C/ ---------- (2) Equations (1) and (2) give the relation between energy of radiation and its frequency or wavelength. A radiation which has higher the frequency or lowers the wavelength has more energy. For example, violet light is larger frequency has more energy than red light which is having lower frequency. Atomic Structure & Basics of Chemistry When the gases or vapor of chemical substances are analyzed with heated by electric spark, light is emitted. The color of light depends upon the substance under investigation. Example, sodium gives a yellow light and potassium produces violet color. This type of spectrum consists of sharp well- defined line each corresponding to a definite frequency (or wavelength). Such spectrum is called line spectrum of discontinuous spectrum. The line spectrum also known as atomic spectrum because, it is obtained from atoms by the application of heat or other energy (electric). Each element gives a unique spectrum irrespective of even the form in which it present. Example, Sodium always gives lines at 589nm and 589.6nm (yellow). From this reason line spectra are also called as finger prints of atoms. 2.7.3 Absorption Spectra When a continuous electromagnetic radiation (white light) is allowed to pass through a gas or a solution of some salt and the transmitted light is analyzed. We obtain a spectrum in which dark lines are observed. These dark lines indicate that the radiations of corresponding wavelength have been absorbed by the substances from the white light. Such a spectrum containing few dark lines due to absorption of light is known as absorption spectrum. The dark lines of wavelengths are also character of a substance. Example, absorption of Na consists of lines at 589nm and 589.6nm same as that of emission spectrum. 2.7.4 Emission Spectrum of hydrogen atom The spectrum of hydrogen atom has played a very important role in the development of atomic structure. The spectrum of hydrogen atom can be obtained by passing an electric discharge through the hydrogen gas taken in the discharge tube under low pressure. The emitted light (radiation) is analyzed by spectroscope. The spectrum consists of large number of lines appearing in different wavelengths (visible, UV, IR regions). Atomic Structure & Basics of Chemistry When a photon is emitted, ni > nf. Consequently the term is negative and E is positive. If when energy is absorbed, ni < nf and the term is positive, so E is positive. The brightness of a spectral line depends many photons of the same wavelength are In equation (3) by substituting RH=2.18x10-18J, h=6.623x10-34 J.sec. and C=3x108m/sec, we found that RH/hC is 1.097x107 m, which is the constant given in the Balmer equation. The emission spectrum of hydrogen includes a wide range of wavelengths from the IR to UV, and lists the series of transitions in the hydrogen spectrum; they are named as, Series nf ni Spectral Lyman 1 2,3,4,…… Balmer 2 3,4,5,…… Visible Paschen 3 4,5,6,……. IR Brackett 4 5,6,7,……. IR Pfund 5 6,7,8,……. IR The above equation (3), and lines only applicable for hydrogen atom. But other atoms lines are complicated and equation is also not obeyed. The energy levels of hydrogen atom with series of emission spectral transition are labeled with their principal quantum numbers. The spectrum of hydrogen atom holds the key to the inner structure of atom. Bohr proposed the theory of atom based on Planck’s quantum theory and spectra of hydrogen atom. 2.8 Bohr model of the atom 1. Maxwell’s law is not applicable to electron in an atom. 2. An atom consists of a massive positively charged nucleus. The electrons revolving around the nucleus in a certain circular orbits without radiation energy. These non-radiating orbits are known as stationary states. 3. Each orbit is at different distance from the nucleus and they are associated with definite Atomic Structure & Basics of Chemistry 3. Because of electron transition between two stationary states of E1 and E2, then the energy is emitted or absorbed in the form of quanta. The frequency, of the emitted or absorbed radiation is given by, h = E2 – E1 4. Since energy cannot be lost continuously, and an electron continuously to move in a particular energy level without losing energy. Such a state of the atom is known as ground state. On gaining energy from an external source electron jumps from a lower energy level to a higher energy level is known as excited state. However the excited state is unstable, and excited electron jumps down from higher to lower energy level (either directly or in steps) by losing energy in the form of electromagnetic radiation. This accounts for the spectral lines in the hydrogen spectrum. 2.8.1 Limitations of Bohr’s Theory: 1. It is applicable only to hydrogen atom or hydrogen like ions which contain only one electron. 2. The spectra of multi electron system cannot be explained by Bohr’s theory. 3. The experimental value of ionization energy and the value calculated from Bohr’s theory does not agree. 4. Bohr Theory cannot explain the mode of formation of bonds between atoms. 5. It gives the flat model of the orbit. 6. By using spectroscope of high resolving it is observed that each line in the spectrum is split into a number of component lines, differing slight in their frequencies, this is called fine structure of spectral lines. Bohr’s theory fails to explain the fine spectrum. 7. Bohr’s theory predicts definite orbits for electrons. This goes against the modern ideas of wave nature of electron. 2.9 Dual Nature of the Electron De Broglie stated that if light waves can behave like a stream of particles (photons), then particles such as electrons could possess wave properties. He argued that if an electron does behave standing wave in the hydrogen atom, the Atomic Structure & Basics of Chemistry Applying Heisenberg uncertainty principle to the H atom, we see that in reality the electron does not in orbit the nucleus in a well-defined path, as Bohr thought. If it did, we could determine precisely the position of the electron and its momentum at the same time, a violation of the uncertainty principle. Significance: 1. The principle is no significance for macro objects. Because x p for ion ball having mass 100. Example, x v= h/4m = 6.623x10-27 erg.sec / 4x3.142x100 = 0.526x10-29 cm2sec-1 2. The principle cannot be neglected for microscopic particle like electron. Example, x v for electron having mass(m) is 9.1x10-28g x v = h/4m = 6.62x10-27 erg.sec ------------------------------- 4x3.142x9.1x10-28 kg It is applicable and cannot be ignored. Uncertainty principle and idea of probability picture of an atom. According to bohr’s model, electrons revolved in well defined circular orbits. But uncertainty principle is not possible to know the exact position and path of electron. This new approach called wave mechanics. 2.10.2 Schrodinger wave mechanics In 1927 Erwin Schrodinger (Austrian physicist) used a complicated mathematical technique formulated an equation that describes the behavior and energies of submicroscopic particles. In general the Schrodinger equation requires advanced calculus to solve and it is important to know that the equation incorporates both particles behavior in terms of mass, m and wave behavior in terms of wave function  (psi), which depends on the location in space of the system. According to wave theory, the intensity of light is proportional to the square of the amplitude of the wave or 2. The most likely to find a photon is where the intensity is greatest, that is, where the value of 2 is greatest. Schrodinger wave equation Atomic Structure & Basics of Chemistry quantum number (n), the angular momentum quantum number (l) and the magnetic quantum number (ml). These quantum numbers will be used to describe atomic orbitals and label the electrons. A fourth quantum number, spin quantum number describes the behavior of specific electron. 2.12.1 Principal quantum number (n) The principal quantum number (n) can have integer values 1,2,3,……… The value of n determines the energy of an orbital. In the case of hydrogen atom or single electron atomic ions, such as Li2+ and Be3+, the energy of these determined by their principal quantum number. For other atoms, the energy also depends on the angular momentum quantum number (l). The size of an atom also depends on the value of n. The orbitals of the same quantum state “n” are said to belong to the same shell. Shells are designates by the following letters, Letter K L M N O P n 1 2 3 4 5 6 2.12.2 Angular momentum or Azimuthal quantum number This quantum number distinguishes the orbitals of a given “n” having different shapes. The values of l depends on the value of principal quantum number, it can have any integer value from 0 to n- 1. With each shell of quantum number, n, there are n different kinds of orbitals, each with a different shape denoted by an “l” quantum number. For example, If, n = 1 has only one l i.e., l=0 n = 2 has two l values i.e., l=0,1 n = 3 has three l values i.e., l=0,1,2 n = 4 has four l values i.e., l=0,1,2,3 An electron has a n value is 3, the possible value for l are 0,1 and 2. Thus, with in the M shell (n=3), there are three kinds of orbitals, each having different shape. The energy also depends on the “l” quantum number, orbitals of the same ‘n’ but different, ‘l’ values are said to belong different sub-shells of a given shell, usually denoted by letters as, Atomic Structure & Basics of Chemistry 2.12.4 Spin quantum number According to electromagnetic theory, a spinning charge generates a magnetic field, and motion that causes an electron to behave like a magnet. The following figure shows the two possible spinnin stats, and author profiles for this publication at: https://www.researchgate.net/publication/326460294 of Chemistry READS 162,043 are also working on these related projects: for Processing of Fish Skin into Leather View project sugar technolgy researches View project by Ramesh Duraisamy on 18 July 2018. the downloaded file. of Chemistry Table of Contents 2.4.2 Characteristic of electromagnetic radiation....................................................... 15 2.4.3 Properties of Waves and Electromagnetic radiation....................................................... 15.......................................................................4 2.4.4 Characteristics of wave motion............. 15 ....................4 2.5 Electromagnetic Radiation.......................... 16 2.6 Quantum effects and photons...................... 16 2.6.1 Planck’s Quantization Energy.............. 17 ..............5 2.6.2 Photoelectric effect.............................. 17 .............5 2.7 Spectrum.................................................... 17 ......................5 2.7.1 Atomic Spectra.................................... 17 2.7.2 Emission Spectra.................................. 17 .......................5 2.7.3 Absorption Spectra............................... 18 2.8 Bohr model of the atom.............................. 19 ...................7 2.8.1 Limitations of Bohr’s Theory:.............. 20 ..............7 2.9 Dual Nature of the Electron........................ 20 ...............7 2.10 Quantum mechanics.................................. 20 ...............8 2.10.1 Heisenberg Uncertainty Principle....... 20 2.10.2 Schrodinger wave mechanics............. 21 2.11 The Quantum mechanical Description of the hydrogen atom.................................................. 21 2.12 Quantum Numbers.................................... 21 10 2.12.1 Principal quantum number (n)............ 22 ................... 10 2.12.3 Magnetic quantum number................. 22 ................. 10 2.12.4 Spin quantum number........................ 23 2.13 Atomic Orbitals........................................ 23 2.13.1 The energies of Orbitals..................... 23 ............... 12 2.13.2 Electron Configuration....................... 24 2.13.3 General Rules for Assigning Electrons .................. 12 to Atomic Orbitals........................................ 24 12 2.13.4 Method of writing electronic ................. 12 configuration:............................................... 24 13 nlx method and Box method.......................... 24 ................ 14 2.14 Pauli Exclusion Principle.......................... 25 2.15 Hund’s Rule.............................................. 25 2.15.1 Paramagnetism and Diamagnatism..... 25......................................................................... 14 2.15.2 The shielding effect in Many-Electron 14 Atoms........................................................... 25 2.16 The Building-Up Principle........................ 26 Properties............. 28 5.7 Precipitation Reactions............................... 53 28 5.7.1 Solubility............................................. 53 5.8 Molecular equations and Ionic equations..... 55 ............... 28 5.9 Acid-base neutralization.............................. 57 28 5.10 Oxidation-Reduction Reactions................. 57 5.10.1 Oxidation Number.............................. 58 ..................... 28 5.10.2 Types of redox reactions.................... 60 5.10.3 Balancing of Redox reactions............. 62 ................ 29 ................ 30 elements..... 32 metals)....... 32 ................. 33 ................... 33 in the p-block ....................... 33 34 ................ 35 family)..................................................................... 36 (Nitrogen ........................ 37 group VIA)..................................................................... 38 metals)......................................................................... 39 Geometry........ 40 40 ................ 42 43 44 44 44 ................ 44 46 47 ................ 50 .................. 51 Page 3  Take notes and review them immediately after class.  Use the five R's of note taking 1. Record: The meaningful ideas and concepts 2. Reduce: After class, summarize the main ideas and concepts 3. Recite: Say out loud in your own words the main ideas of the class. everyone is 4. Reflect: Take a few minutes to ponder slightly over the main ideas of the class. different set of 5. Review: Once a week, review the ideas of bring to the all lectures. 1.1.2 Studying  Study before each class. The material will be fresh in your mind.  Allow time for sleep.  Set realistic goals for yourself. You can reward yourself for being successful. If you waste all process afternoon and then set a goal of studying chemistry from 6:00 p.m. to midnight, you may or her mind, not be realistic. 1.1.3 Develop an effective set of study skills  Pace your studying throughout the week and throughout the semester  Get enough rest and otherwise take care of your health.  Find a place in which you can study effectively  Form a study group with classmates and discuss words, concepts, problem-solving strategies, etc. to the benefit of all the members of the group. 1.1.4 Develop your problem-solving skills  Sharpen your basic mathematical skills  Learn how to use your calculator effectively.  Learn & practice using units for physical quantities.  If you need a piece of information, look it up using the index of your text or other suitable source. Don't be afraid to be wrong, as Sir Isaac Newton said many years ago, "Truth comes more easily out of error than out of confusion." of lecture Page 4 be determined directly. Microscopic properties, on the atomic or molecular scale, must be determined by an indirect method. A measured quantity is usually written as a number with an appropriate unit. The same is true in chemistry; units are essential to stating measurements correctly. each of which 1.3.1 SI units Many years’ scientists recorded measurements in metric units, which are related decimally (i.e. by powers of 10). On 1960‘s the general conference of weights and measures notified, the international system of units (SI). The SI system consists of the following seven units, from which other units of weight and measure can be derived. Measurements that we will utilize frequently in our study of chemistry include time, mass, volume, density and temperature. SI base units: --------------------------------------------------------------- Physical property Name of unit Symbol --------------------------------------------------------------- Length meter m Mass kilogram kg is called a Time second s Electric current Ampere A Temperature Kelvin K Luminous intensity Candela cd Amount of substance mole mol --------------------------------------------------------------- the powers to which 10 is raised as, common prefixes used in the metric system --------------------------------------------------------------- Prefix Symbol Factor Example --------------------------------------------------------------- pico p 10-12 1 picoeter (pm) -9 nano n 10 1nano gram (ng) used in micro  10-6 1microliter (L) . milli m 10-3 2milliseconds(ms) centi c 10 -2 5centimeters(cm) deci d 10-1 1 deciliter (dL) 3 kilo k 10 1kilometer(km) mega M 106 3 mega grams (Mg) 9 giga G 10 5 giga meters (Gm) provide tera T 1012 1 tera litre (TL) --------------------------------------------------------------- Page 5 1 g/cm3 = 1g/mL = 1000 kg/m3 1g/L = 0.001g/mL The term specific gravity denotes the ratio of the density of a substance to the density of a reference substance. Usually, the reference substance for solids and liquids is water. Density of substance Specific gravity = ----------------------------------------- Density of reference substance Specific gravity, being the ratio of two densities, has no units. Usually the specific gravity of any substances is numerical. Workout.1 Calculate the density and specific of a body gravity of a body that has a mass of 321g and a an accident volume of that 45 cm3 at 250C. term weight is Mass 321g Density = --------------- = ------------ = 7.13 g/cm3 Volume 45 cm3 Density of sample Specific gravity = -------------------------- dium Density of water The density of water at 250C is 0.99 g/cm3, specific gravity of sample is 7.13. Workout.2 What is the mass in kg of 10.5g (39.7L) of gasoline with a specific gravity of 0.82? Because gasoline has a specific gravity of 0.82 the density of 1mL of gasoline is 0.82 times that of 1mL of water. Density = Specific gravity x density of water = 0.82 x 0.997 g/cm3 D ------- = 0.82 g/cm3 (or 0.82 g/mL) Conversion of 39.7 L to mL followed by multiplication by the mass of 1mL of gasoline gives the total mass. 1000 mL Volume = 39.7L x ------------- = 39700 mL 1L Mass = Density x Volume = 0.82 g/mL x 39700 mL = 33,000 g Mass = 33kg Page 6 carbon atom. Because we have 3x1022 atoms of carbon we can make, 1 CCl4 molecule 22 3 x 10 C atoms x ------------------------- 1 C atom 22 = 3x10 CCl4 molecules and the SI Problem.5 How many moles of chlorine atom are required to react with 0.050 mol of carbon atoms to convert all of the carbon atoms to CCl4. The formula CCl4 indicates that there are 4 Cl atoms for each C atom in a molecule, thus we must take four times as many Cl atoms as C atoms, 6.023x1023 C atoms 0.050 mol C atoms x ------------------------------- 1 mol C atoms 22 = 3x10 C atom 4 Cl atoms 3x10 C atoms x ---------------- = 1.2x1023 Cl atoms 22 1 C atoms calculate the number of moles of Cl atoms from the definition of a mole, 1 mole of Cl atoms contains 6.023x1023 Cl atoms [1 mol Cl atom / 6.023x1023 Cl atoms) 1 mol Cl atom 23 1.2x10 Cl atoms x ----------------------------- 6.023x1023 Cl atoms = 0.20 mol of Cl atom 1.4 The Classification of Matter The chemistry deals the study of matter and the changes it undergoes. Matter is anything that occupies space and has mass. Thus, everything in the universe has a “chemical” connection. Chemical distinguish among several subcategories of matter based on composition and properties. The classification of matter includes substances, mixtures, elements and compounds, as well as atoms and molecules. ------------------ 1.4.1 Three states of matter All substances, at least in principle, can exist in three states: solid, liquid, and gas. The gases differ of carbon. from liquids and solids in the distances between the for each molecules. In a solid, molecules are held close Page 7 All measurable properties of matter fall into one of two additional categories: extensive properties and intensive properties. The measured value of an extensive property depends on how much matter is being considered. Mass, which is the quantity of matter in a given sample of a substance, is an extensive property. More matter means more mass. Values of the same extensive property can be added together. For example, two copper pennies will have a combined mass that is the sum of the masses of each penny, and the length of two tennis courts is the sum of the lengths of each tennis court. Volume, defined, as length of cubed, is another extensive property. The value of an extensive quantity depends on the amount of matter. The measured value of an intensive property does not depend on how much matter is being considered. Density, defined as the mass of an object divided by its volume, is an intensive property. We consider the temperature. Suppose that we have two beakers of water at the same temperature. If we combine them to make a single quantity of water in a larger beaker, the temperature of the larger quantity of water will be the same as it was in two separate beakers. Unlike mass, length, and volume, temperature and For other intensive properties are not in additive. 1.5.2 Atoms and Molecules of ice by An atom is the smallest particle of an element that temperature can enter into a chemical combination. The word Water differs atom comes from the Greek word atomos, which so means indivisible. John Dalton, an English chemist the water to and physicist, presented an atomic theory and melting point supported with quantitative measurements, found it. Similarly, A molecule is the smallest particle of an element or than air, we compound that can have a stable, independent existence. A molecule may consist of a single atom, gas as in helium, of two or more identical atoms, as in describes a nitrogen and sulfur, or of two or more different in order to atoms, as in water. Water has a definite composition a chemical and a set of chemical properties that enable us to change, the recognize it as a distinct substance. Each water gas, will molecule is a unit that contains two hydrogen atoms is a different and one oxygen atom. Subdivision of a water we cannot molecule results in the formation of the gases means of a hydrogen and oxygen, each of which has properties Page 8 For convenience, chemists use symbols of one or two letters to represent the elements. The first letter of a symbol is always capitalized, but any following letters are not. For example, Co is the symbol for the element cobalt, whereas CO is the formula for the carbon monoxide molecule. Names and symbols of some of the more common elements are presented as, ----------------------------------------------------------------- Name Symbol Name Symbol ----------------------------------------------------------------- Aluminum Al Fluorine F Oxygen O Arsenic As Gold Au Phosphorus P Barium Ba Hydrogen H Platinum Pt Bismuth Bi Iodine I Potassium K Bromine Br Iron Fe Silicon Si Calcium Ca Lead Pb Silver Ag Carbon C Magnesium Mg Sodium Na Chlorine Cl or Manganese Mn Sulfur S Chromium Cr Mercury Hg Tin Sn Cobalt Co Thus, Nickel Ni Tungsten W Copper Cu Nitrogen N to dryness. Zinc Zn the water --------------------------------------------------------------- The symbols of some elements are derived from their Latin names-for example, Au from aurum (gold), Fe from ferrum (iron), and Na from natrium (sodium), while most of them come from English names. Atoms of most elements can interact with one another to form compounds. For example, hydrogen gas burns in oxygen gas to form water, which has properties that are distinctly different from those of the starting materials. Water is made up of two parts hydrogen and one part oxygen. Thus, water is a compound, a substance composed of atoms of two or more elements chemically united in fixed proportions. Page 9 atom (Cl) can gain an electron to become the chloride ion Cl- Cl atom: 17 protons and 17 electrons Cl- ion : 17 protons and 18 electrons Sodium chloride (NaCl), table salt, is an ionic compound because it is formed from cations and anions. An atom can lose or gain more than one electron. The ions Na+ and Cl- are called monatomic ions because they contain only one atom. In addition, two or more atoms can combine to form an ion that has a net positive or net negative charge. Examples of ions formed by the loss or gain more than one electron are Mg2+,Fe3+,S2- and N3-. Polyatomic ions such as OH- (hydroxide ion), CN- (cyanide ion) and NH4+ (ammonium ion) ions are containing more than one atom. compound 1.6 Significant Figures Except when all the numbers involved are integers, it is often impossible to obtain the exact value of the quantity under investigation. For this reason, it is important to indicate the margin of error in a measurement by clearly indicating the number of significant figures, which are the meaningful digits in a measured or calculated quantity. When significant figures are used, the last digit is understood to be uncertain. 1.6.1 Guidelines for using Significant figures We must always be careful in scientific work to write the proper number of significant figures. In general, it is fairly easy to determine how many significant figures a number has by following these rules:  Any digit that is not zero is significant. Thus 815 cm has three significant figures, 1.374 kg has four significant figures, and so on.  Zeros between nonzero digits are significant. Thus 608 m contains three significant figures, 30,403 kg contains five significant figures, and so on.  Zeros to the left of the first nonzero digit are not significant. Their purpose is to indicate the placement of the decimal point. For example, 0.08 L contains one significant figure; 0.000756 g contains three significant figures, and so on. A chlorine Page 10 only two digits after the decimal point. If the first digit following the point of rounding off is equal to or greater than 5, we add 1 to the proceeding digit. Thus 8.727 rounds off to 8.73 and 0.425 rounds of to 0.43.  In multiplication and division, the original number that has the smallest number of significant figures determines the number of significant figures in the final product or quotient. The following examples illustrate this rule: 2.8 x 4.5039 = 12.61092 --- round off to 13 6.85/112.04 = 0.0611388789 round off to 0.0611  Keep in mind that exact numbers obtained from definitions or by counting numbers of objects can be considered to have an infinite number of significant figures. If an object has a mass of 0.2786g, then the mass of eight such object is, 0.2786 g x 8 = 2.229 g two significant We do not round off this product to one significant figure, because the number 8 is 8.00000 …… by definition. Similarly, to take the average of the two measured lengths 6.64 cm and 6.68 cm, we write 6.64 cm + 6.68 cm ----------------------- = 6.66 cm 2 Because the number 2 is 2.00000 …….. by definition. cannot decimal point Consider off to 90.4 off to 1.98 we simply first of them is 8.72 if we want Page 11 the right kinds of elements, but specific numbers of these atoms as well. This idea is an extension of law published by Joseph Proust (in 1799), a French chemist. Proust’s law of definite proportions states that different samples of the same compound always contain its constituent elements in the same proportion by mass. Dalton’s 2nd hypothesis supports another important law, the law of multiple proportions. According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. Dalton’s theory explains the laws of multiple proportions, different compounds made up of the same elements differ in the number of atoms of each kind that combine. For example, carbon forms two stable compounds with oxygen, namely carbon monoxide and carbon dioxide. Third hypothesis It is another way of stating the law of conservation of mass, which is that matter can be neither created nor destroyed. Since matter is made of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well. present is 2.2 The structure of the Atom On the basis of Dalton’s atomic theory, we can define an atom as the basic unit of an element that can enter into chemical combination. Dalton imagined an atom that was both extremely small and indivisible. However, a series of investigations, clearly demonstrated that, atoms actually posses internal structures; that is, they are made up of even smaller particles, which are called subatomic particles. This research discovered the three such particles, named as electrons, protons and neutrons. 2.2.1 The Electron One device used to investigate the atomic structure was a cathode ray tube; it is a glass tube from which most of the air has been evacuated. When the two metal plates are connected to a high-voltage source, the negatively charged plate, called the cathode, emits a ray. The cathode ray is drawn to the positively charged plate, called anode, where it Page 12 emit very unusual rays. This highly energetic radiation penetrated matter, darkened covered photographic plates, and caused a variety of substances to fluoresce. Since a magnet could not deflect these rays, they could not contain charged particles as cathode rays. Rontgen called them X rays because their nature was not known. After Rontgen’s discovery, Antoine Becquerel, began to study the fluorescent properties of substances, he found that exposing thickly wrapped photographic plates to a certain uranium compound caused them to darken, even without the stimulation of cathode rays. Like X rays, the rays from the uranium compound were highly energetic and could not be deflected by a magnet, but they differ from X rays because they arose spontaneously. One of Becquerel’s students, Marie Curie, suggested the name radioactivity to describe this spontaneous emission of particles and/or radiation. Since then, any element emits the radiation as spontaneously is said to be radioactive. Three types of rays are produced by the decay, or breakdown, of radioactive substances such as uranium. Two of the three are deflected by oppositely charged metal plates (shown in figure). Alpha () rays consist of positively charged particles, called  particles, and therefore are deflected by the positively charged plate. Beta () rays, or  particles, are electrons and are deflected by the negatively charged plate. The third type of radioactive radiation consists of high-energy rays called gamma () rays. Like X rays,  rays have no charge and are not affected by an external field. 2.2.2 The proton and Nucleus By the early 1900s, two features of atoms had become clear: they contain electrons, and they are electrically neutral. To maintain electric neutrality, an atom must contain an equal number of positive and negative charges. Therefore, Thomson proposed that an atom could be thought of as a uniform, positive sphere of matter in atoms. The New Zealand physicist Ernest Rutherford, studied with Thomson at Cambridge University, decided to use  particles to probe the structure of (in atoms. Rutherford carried out a series of and metals to Page 13 equal to the number of electrons, so the atomic number also indicates the number of electrons present in the atom. For example, the atomic number of nitrogen is 7. This means that each neutral nitrogen atom has 7 protons and 7 electrons. Mass number The mass number (A) is the total number of protons and neutrons present in the nucleus of an atom of an element. Except for the most common form of hydrogen, which has one proton and no neutron, all atomic nuclei contain both protons and neutrons. In general the mass number is given by, Mass number = No. of protons + No. of neutrons = Atomic number + No. of neutrons The number of neutrons in an atom is equal to the difference between the mass number and the atomic number, (A-Z). For example, the mass number of fluorine is 19 and the atomic number is 9 (indicate 9 protons in the nucleus). Thus the number of neutrons in an atom of fluorine is 19-9=10. The atomic number, number of neutrons and mass number all must be positive integer (whole numbers). The accepted way to denote the atomic number and mass number of an atom of an element is represented as ZXA, A is mass number and Z is an atomic number. Isotopes Most of the elements have two or more isotopes; atoms that have the same atomic number but different mass numbers. For example, there are three isotopes of hydrogen. One, simply known as hydrogen, has one proton and no neutrons. The deuterium isotope contains one proton and one neutron, and tritium has one proton and two neutrons. Thus, for the isotope of hydrogen, we write 1H1 (hydrogen), 1H2 (deuterium) and 1H3 (tritium). As another example, consider two common isotopes of uranium with mass numbers of 235 and 238, respectively: 92U235, 92U238. 2.4 Nature of light and Electromagnetic radiation 2.4.1 Corpuscular Theory Earliest view of light due to Newton regarded light (Z) is the as made up of particles (corpuscles). This is due to atom of an of protons is Page 14 forms of electromagnetic radiation. The point of maximum upward displacement is called crest, and maximum downward displacement is called trough. Thus, waves may be considered as a continuous sequence of alternating crest and troughs. 2.4.4 Characteristics of wave motion A wave was characterized by its wavelength and frequency. 1. Wave length () The wavelength denoted by  (lambda) is defined as the distance between any two adjacent peaks (crests) or troughs on successive waves. Unit – A0 (Angstron) 1A0 = 10-8 cm = 10-10 m SI unit:  (micrometer), m (milli micron) , nm (nanometer), pm (picometer). 1A0 = 10-10m; 1 = 10-6m; 1m = 10-9m 1nm = 10-9m; 1pm=10-12m reaches 2. Frequency () The frequency, (nu) is the number of waves or cycles per second pass through a given point. In following figure shows the frequency corresponds to the number of times per second that moves through a complete cycle of upward and downward motion. The frequency is expressed in the unit of cycles per second (s-1) or Hertz (Hz). 1 Hz = 1 cycles per second = C/ 3. Amplitude: It is the vertical distance from the midline of a wave to the peak or trough. Height of crest or depth of trough i.e. intensity or brightness. 4. Velocity (C): Distance travelled by a light in one second. The unit of velocity is ms-1 or cms-1 5. Wave number (): The number of waves per unit length. Unit: cm-1 or m-1, = 1/ of quiet Relationship: C =  x ; = C/ 1/  = ; = C Exercise.1 Calculate the speed of a wave whose wavelength and frequency are 17.4cm and 87.4 Hz, respectively. This is directly solved by the application of c=  c= 17.4 cm x 87.4 Hz ; c = 17.4cm x 87.4/s = 1.52 x 103 cm/s Page 15 planes, and his model describes how energy in the form of radiation can be propagated through space as vibrating electric and magnetic fields. Electromagnetic radiation is the emission and transmission of energy in the form of electromagnetic waves. r 10-6 10-4 10-3 10-1 Far Near Vis Near Far IR Microwave Radio IR Radar TV, FM 1015 1014 1012 1010 Types of electromagnetic radiation Red 600nm 700nm ranges from 400nm (violet) to 700nm (red) than 800nm and ultraviolet radiation has wavelengths less than 400nm. 2.6 Quantum effects and photons Issac Newton, who studied the properties of light in the 17th century, believed that light consisted of a beam of particles. In 1801, British physicist Thomas Young showed that light like waves could be diffracted. By the early part of the twentieth century, the wave theory of light appeared to be well entrenched. But in 1905, the German physicist Albert Einstein discovered that he could explain a phenomenon known as the photoelectric effect by postulating that light had both wave and particle properties. Einstein based this idea on the work of the German physicist Max Planck. Page 16 2.6.2 Photoelectric effect Einstein used this photon concept to explain the photoelectric effect. The photoelectric effect is the ejection of electrons from the surface of a metal or from another material when light shines on it. Electrons are ejected, only when the frequency of light exceeds at a certain threshold value of the particular metal. For example, violet light will cause potassium metal to eject electrons; not in red light (which has lower frequency). When the photon hits the metal, its energy h is takes up by the electron. The photon cease to exist as a particle; it is said to be absorbed and emits the electron as radiations (light). The wave and particle pictures of light should be regarded as complementary views, called wave- particle duality of light. The equation E = h displays this duality. i.e., E is the energy of a light particles or photon and is the frequency of the associated wave. 2.7 Spectrum When white light from Sun is passed through a prism, it split into a series of color band known as rainbow colors (VIBGYOR). This means that Sunlight is composed of collection of electromagnetic waves having different wave length. The splitting of light into seven colors is known as dispersion and the series of color bands is called a spectrum. In this spectrum, there is continuity of colors such a spectrum is known as continuous spectrum. 2.7.1 Atomic Spectra Unlike the spectrum obtained by analyzing the sunlight, the spectra of atoms of elements can be made to emit energy by subjecting them to electric discharge or by heating are not continuous (discontinuous). The spectrum of atoms consists of sharp well-defined lines or bands corresponding to definite frequencies. There are two types of atomic spectra. (i) Emission spectra (ii) Absorption spectra. 2.7.2 Emission Spectra Emission spectra are obtained when the radiations emitted from substances are analyzed with the help of spectroscope. Page 17 In 1885, J.J.Balmer developed a relationship among the different wavelengths,  in the visible region in hydrogen spectrum, could be reproduced by simple formula: i.e., 1/ = (cm-1) = 109678 (1/22 – 1/ni2) or (cm-1) = 1.097 x 107(1/22 – 1/ni2) ni is an integer but greater than 2, ni = 3,4,5,….. These series of lines which appear in visible region are named as Balmer series. Afterwards, a series of spectral lines of H2 atom in different regions were discovered. These series of lines named after the name of its discoverers. These are Lyman, Balmer, Paschen, Brackett and Pfund series. Lyman – UV region ; Balmer – Visible Paschen, Brackett and Pfund – IR region As the other series of hydrogen spectral lines were discovered, a more general expression was found by Rydberg, and is known as Rydberg equation derived as, E = - RH / n2 ------------ (1) n = 1,2,3,..,  (for hydrogen atom) Where, RH is Rydberg constant has the value 2.18x10-18J n is principal quantum number negative sign indicates arbitrary convention. The energy levels of the electron in the hydrogen atom do undergo a transition of electron between energy levels; electron loses energy, which is emitted as photon. The difference between the energies of the initial and final state is: E = Ef – Ei or Ei - Ef From equation (1), Ef = - RH(1/nf2) and Ei = - RH (1/ni2) Therefore, E = - RH / nf2 – (- RH/ni2) = RH/ni2 – RH/nf2 E = RH [1/ni2 – 1/nf2] --------- (2) Because this transition results in the emission of a current photon of frequency, and energy h, we can write: E = h = RH [1/ni2 – 1/nf2] and, we have = c/, we can write this hC/ = RH [1/ni2 – 1/nf2] 1/ = = RH/hC [1/ni2 – 1/nf2] ------- (3) Page 18 amount of energy. Hence, stationary states also called as energy levels. The energy associated with different energy levels increases with increase in distance from the nucleus. The letters of orbit (shell) : K,L,M,N,…… on how The numbers of orbit (shell) : 1,2,3,4,……… emitted. are used to designate the energy levels. 4. Energy associated with an energy level is given by the relation: - 2π2k2Z2me4 En = ------------------------- n2h2 Where, Z – atomic number m – mass of electron h – Planck’s constant e –charge of electron n – orbit number, and k = 9 x 109 Nm2c-2 region k = 1/ 4 π0 ; Ultraviolet 0 being permittivity of empty space and UV For hydrogen atom, Z = 1; - 2π2k2me4 En = ------------------------- n2h2 substituting the value of m, e, π, k and h: EH = - 1312 / n2 kJ/mol For He+, Z = 2  EHe+ = 4 EH For Li2+, Z = 3  ELi+ = 9 EH 1. Different energy levels are not equally spaced. i.e., the energy difference between two successive energy levels is not same. It goes on decreasing with increasing the value of n. 2. Angular momentum of an electron is a whole number multiple of h/2π. i.e. mvr = nh / 2π h - Planck’s constant; m- mass of electron v - tangential velocity; r - radius of the orbit and, r = nh/2πmv The velocity of electron in nth orbit of an atom is given by vn = 2πe2Z / nh, and r = n2h2 / 4π2mZe2 substituting the value of h, π, m, Z and e for hydrogen, Z=1 comes to be 0.53A0 Page 19 wavelength of the wave must fit exactly with circumferences of the orbit. The relation between the circumference of an allowed orbit (2r) and the wavelength () of the electron is given by, 2r = n ---------------- (1) Where, r is the radius of the orbit,  is the wavelength of the electron wave and n=1,2,3, ….. So, energy of electron depends on size of the orbit, its value must be quantized. Thus, the waves can behave like particles and particles can exhibit wavelike properties. So, De Broglie deduced that the particle and wave properties are related by the expression,  = h/mu ----------------- (2) Where, m and u are the wave length associated with a moving particle. Note that the left side of equation (2) involves the wavelike property of wavelength, whereas the right side makes references to mass, a distinct property of particle. Workout: 1. Calculate the de-Broglie wavelength of a ball of mass 0.1kg moving with a velocity 100ms- (Ans: 6.626x10-35m). 2. The mass of an electron is 9.1x10-31kg. The velocity is 800 ms-1. Calculate its wavelength. 2.10 Quantum mechanics Bohr’s approach did not account for the emission spectra of atoms containing more than one electron, such as atoms of He and Li. He did not explain why power, extra lines appear in the hydrogen emission spectrum ordinary when a magnetic field is applied and could not define the precise location of a wave. Describe the problem of trying to locate a subatomic particle that behaves like a wave, Heisenberg formulated the uncertainty principle. 2.10.1 Heisenberg Uncertainty Principle It stated that it is impossible to know simultaneously both the momentum (p) and the position (x) of a particle with certainty. It is stated mathematically, x p  h/4 ---------------------- (3) Where x and p are the uncertainties in measuring the position and momentum. like a Page 20 (E=H) began a new area in physics and chemistry called the quantum mechanics (wave mechanics). 2.11 The Quantum mechanical Description of the hydrogen atom The Schrodinger equation specifies the possible energy states of the electron can occupy in a hydrogen atom and identifies the corresponding wave functions (). These energy states and wave functions are characterized by a set of quantum numbers, with which we can construct a model of the hydrogen atom. In quantum mechanics, concept of electron density gives the probability that an electron will be found in a particular region describes the atom. The square of the wave functions, 2, defines the distribution of electron density in three-dimensional space around the nucleus. According to quantum mechanics, an atomic orbital can be thought as the wave function of an electron in an atom. When we say that an electron is in a certain orbital, it means that the distribution of the electron density or the probability of finding the electron in space is described by the square of the wave function (2) associated with an orbital. Therefore, an atomic orbital has a characteristic energy, as well as a characteristic distribution of electron density. “An orbital is defined as a region in space around the nucleus where the probability of finding the electron is maximum”. The Schrodinger equation is applicable nicely for the simple hydrogen atom with its one proton and one electron, but it turns out that it cannot be solved exactly for many-electron atoms. So, chemists and physicist have learned to get the solution for this kind by applying approximation methods. So, they concluded that the behavior of electrons in many- electron atoms is not the same as in the hydrogen atom. 2.12 Quantum Numbers In quantum mechanics, three quantum numbers are required to describe the distribution of electrons in atoms. These numbers are derived from the mathematical solution of the Schrodinger equation for the hydrogen atom. They are called the principal Page 21 Letter s p d f g h l 0 1 2 3 4 5 n = 2, l = 0,1 has two sub shells (i.e. s, p) n = 3, l = 0,1,2 has 3 sub shells (i.e. s, p, d) n = 4, l = 0,1,2,3 has four sub shells (i.e. s, p, d, f) Thus, if l=0, we have an s orbital; if l=1, we have a s and p orbital and so on. We denote a sub-shell with a particular shell, we write the value of n quantum number for the shell, followed by the letter designation for the sub-shell. Thus, n=1, l=0 has one sub shell ---- s and, maximum number of electrons in each sub-shell: s - 2; p - 6 ; d-10 and f - 14. For example, 2p denotes a sub-shell with quantum numbers n = 2 and l = 1. Correlation between n and l: -------------------------------------------------------------- n l sub-shell no.of sub-shell no.of eles. -------------------------------------------------------------- 1 0 1s one 2 2 0, 1 2s, 2p two 8 3 0, 1, 2 3s, 3p, 3d three 18 4 0, 1, 2, 3 4s, 4p, 4d, 4f four 32 -------------------------------------------------------------- 2.12.3 Magnetic quantum number The magnetic quantum number (ml) describes the orientation of the orbital in space. Within a sub-shell, the value of ml depends on the value of the angular momentum quantum number, l. The allowed values are the integers as follows: -l,….. ,0,……+l (or) –l,(-l+1),…,0,---. (+l-1),+l For l=0 (s subshell), the allowed ml quantum number is 0 only; ther is only one orbital in the s sub-shell. For l=1 (p subshell), ml=-1,0 and +1; there are 3 different orbitals in the p sub-shell. The orbitals have the same sub-shell, but it has different orientation in space (as in px, py and pz). if, n=2, ml = -2,-1,0,+1,+2 has five orientation: dxy, dyz, dxz, dx2 - y2, dz2. if, n=3, ml=-3,-2,-1,0,+1,+2,+3 has seven orbital orientations. But the orientation of ‘f’ orbital is difficult to suggest. In addition each orbital of a given sub-shell has the same energy, note there are (2l+1) orbitals in each sub-shell of quantum number. Page 22 it is this

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