Tark Chemistry Lecture Companion PDF

SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah CHAPTER 2 INORGANIC CHEMISTRY 2.1 INTRODUCTION...

SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah CHAPTER 2 INORGANIC CHEMISTRY 2.1 INTRODUCTION 2.1.1 Inorganic Chemistry 2.1.2 Periodic Table – Periods, Groups, Blocks & History 2.2 TRENDS IN PERIODIC PROPERTIES 2.2.1 Trends in Physical Properties a. Atomic Radius b. Ionic Radius c. Ionization Enthalpy d. Electron Gain Enthalpy e. Electronegativity 2.2.2 Trends in Chemical Properties a. Periodicity of Valence or Oxidation States b. Anomalous Properties of Second Period Elements (Diagonal Properties) 2.3 CHEMICAL BONDING 2.3.1 Ionic / Electrovalent Bond 2.3.2 Covalent Bond 2.3.3 Coordinate Covalent Bond 2.3.4 Metallic Bond 2.3.5 Hydrogen Bond a. Intermolecular H – bond b. Intramolecular H – bond 2.4 MOLE CONCEPT 2.4.1 Mole 2.4.2 Molar Mass 2.4.3 Mass Percentage of Element 2.5 THEORY OF SOLUTION 2.5.1 Types of Solution 2.5.2 Concentration of Solution a. Per cent by weight (% w/w) *Proprietary material of SILVER OAK UNIVERSITY Page | 1  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah b. Mole fraction (x) c. Molarity (M) d. Molality (m) e. Normality (N) 2.6 STRUCTURE OF MATERIAL (SOLIDS) 2.6.1 Ionic Solids 2.6.2 Metallic Solids 2.6.3 Molecular Solids 2.6.4 Network Solids 2.7 CRYSTALLIZATION OF INORGANIC SALTS *Proprietary material of SILVER OAK UNIVERSITY Page | 2  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah INORGANIC CHEMISTRY 2.1 Introduction: 2.1.1 Inorganic Chemistry:  Inorganic chemistry deals with synthesis and behaviour of inorganic and organometallic compounds.  It has applications in every aspect of the chemical industry, including catalysis, materials science, pigments, surfactants, coatings, medications, fuels, and agriculture. 2.1.2 Periodic Table – Periods, Groups, Blocks & History:  The periodic table lists all the elements, with information about their atomic weights, chemical symbols, and atomic numbers.  The arrangement of the periodic table leads us to visualize certain trends among the atoms. (Periodic properties)  The vertical columns (groups) of the periodic table are arranged such that all its elements have the same number of valence electrons.  All elements within a certain group thus share similar properties. Glenn T. Seaborg’s work in the middle of the 20th century starting with the discovery of plutonium in 1940, followed by those of all the trans-uranium elements from 94 to 102 led to reconfiguration of the periodic table placing the actinoids below the lanthanoids. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named Seaborgium (Sg) in his honour. The Periodic Table is arguably the most important concept in chemistry, both in principle and in practice. It is the everyday support for students, it suggests new avenues of research to professionals, and it provides a succinct organization of the whole of chemistry. It is a remarkable demonstration of the fact that the chemical elements are not a random cluster of entities but instead display trends and lie together in families. An awareness of the Periodic Table is essential to anyone who wishes to disentangle the world and see how it is built up from the fundamental building blocks of the chemistry, the chemical elements. ~ Glenn T. Seaborg  Period: Horizontal row in the periodic table, which signifies the total number of electron shells in an element’s atom. *Proprietary material of SILVER OAK UNIVERSITY Page | 3  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  Atomic Number: The number, equal to the number of protons in an atom that determines its chemical properties.  Group: Vertical column in the periodic table, which signifies the number of valence shell electrons in an element’s atom. History & Evolution of Periodic Table: Grouped elements into TRIADS (सारे गामापा), based on Johann Dobereiner 1817 atomic mass & similar properties. Law of Octaves – when placed in order of increasing John Newlands 1863 atomic mass, he found that similar properties appeared every 8 elements. Dimitri Mendeleev Placed elements in order of increasing atomic mass. Lothar Meyer 1869 Credited with creating 1st Periodic Table. Gave us (working separately) “periodic law.” Arranged table in order of increasing atomic number, Henry Moseley 1913 based on x-ray experiments. Gave us modern periodic law & table order. Blocks in Periodic Tables: s – block: *Proprietary material of SILVER OAK UNIVERSITY Page | 4  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration belong to the s – block Elements. p – block:  The p – block Elements comprise those belonging to Group 13 to 18 and these together with the s – block Elements are called the Representative Elements or Main Group Elements. d – block:  These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d – block Elements. f – block:  The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce (Z = 58) – Lu (Z = 71) and Actinoids, Th (Z = 90) – Lr (Z = 103). The last electron added to each element is filled in f- orbital. These two series of elements are hence called the Inner – Transition Elements (f – block Elements). Group Names: Group Number Group Name 1 Alkali Metals 2 Alkaline Earth Metals 3 – 12 Transition Metals 15 Pnictogens 16 Chalcogens 17 Halogens 18 Nobel Gases Metals, Non – metals & Metalloids:  Metals: a type of solid substance that is usually hard and shiny and that heat and electricity can travel through *Proprietary material of SILVER OAK UNIVERSITY Page | 5  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  Non – metal: a type of solid substance that is usually not that much hard and shiny and that heat and electricity cannot travel through  Metalloids: an element whose properties are intermediate between those of metals and solid non-metals or semiconductors 2.2 Trends in Periodic Properties: 2.2.1 Trends in Physical Properties: a. Atomic Radius:  The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the centre of the nucleus to the boundary of the surrounding shells of electrons.  Atomic radius increases, as we move down a group.  The number of energy levels increases as we move down a group as the number of electrons increases. Each subsequent energy level is further from the nucleus than the last.  Atomic radius decreases, as we across a period. *Proprietary material of SILVER OAK UNIVERSITY Page | 6  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  As we go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge”. *Proprietary material of SILVER OAK UNIVERSITY Page | 7  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah Atomic Radius 152 160 140 112 120 100 85 Atomic Radius 77 80 70 73 72 60 0 1 2 3 4 5 6 7 8 9 10 b. Ionic Radius:  Ionic radius, rion, is the radius of a monatomic ion in an ionic crystal structure.  Anions (negative ions) are larger than their respective atoms.  Electron-electron repulsion forces them to spread further apart. Electrons outnumber protons; the protons cannot pull the extra electrons as tightly toward the nucleus.  Cations (positive ions) are smaller than their respective atoms.  There is less electron-electron repulsion, so they can come closer together. Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly.  If the electron that is lost is the only valence electron (in case of alkali metals) so that the electron configuration of the cation is like that of a noble gas, then an entire energy level is lost. In this case, the radius of the cation is much smaller than its respective atom. *Proprietary material of SILVER OAK UNIVERSITY Page | 8  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah c. Ionization Enthalpy:  Ionization energy is the minimum amount of energy required to remove the most loosely bound electron of an isolated neutral gaseous atom or molecule.  Second ionization enthalpy is the energy required to remove the second most loosely bound electron.  Energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive.  The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom.  As we move down a group, first ionization energy decreases.  Electrons are further from the nucleus and thus easier to remove the outermost one. Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus. It therefore becomes easier to remove the outer electron Shielding Effect / Screening Effect:  Reduction in the effective nuclear charge on the electron cloud, due to a difference in the attraction forces on the electrons in the atom. *Proprietary material of SILVER OAK UNIVERSITY Page | 9  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  As we move across a period, first ionization energy increases.  As we move across a period, the atomic radius decreases, that is, the atom is smaller. The outer electrons are closer to the nucleus and more strongly attracted to the centre. Therefore, it becomes more difficult to remove the outermost electron. Exceptions to First Ionization Energy Trends: 1. Xs2 > Xp1 e.g. 4Be > 5B The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital. Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital. 2. Xp3 > Xp4 e.g. 7N > 8O After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired. The electron-electron repulsion makes it easier to remove the outermost, paired electron. d. Electron Gain Enthalpy:  Amount of energy released when an electron is added to an isolated gaseous atom.  During the addition of an electron, energy can either be released or absorbed.  The electron gain enthalpy for halogens is highly negative because they can acquire the nearest stable noble gas configuration by accepting an extra electron. *Proprietary material of SILVER OAK UNIVERSITY Page | 10  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  Noble gases have large positive electron gain enthalpy because the extra electron has to be placed in the next higher principal quantum energy levels there by producing highly unstable electronic configuration.  After the addition of 1 electron, the atoms become negatively charged and the second electron is to be added to a negatively charged Ion. But the addition of second electron is opposed by the electrostatic repulsion and hence the energy has to be supplied for the addition of second electron. The second electron gain enthalpy of an element is positive.  As we move down a group, electron gain enthalpy decreases.  As we move down a group, both the atomic size and the nuclear charge increases. But the effect of increase in atomic size is much more pronounced then the nuclear charge. With increase in atomic size, the attraction of the nucleus for the incoming electron decreases and hence the electron gain enthalpy becomes less negative.  As we move across a period, electron affinity increases.  As we move across a period from left to right the atomic size decreases and the nuclear charge increases. Both these factors tend to increase the attraction by the nucleus for the incoming electron and hence electron gain enthalpy becomes more and more negative in a period from left to right. Exceptions:  Among non-metals, the elements in the first period have lower electron affinities than the elements below them in their respective groups.  Elements with electron configurations of Xs2, Xp3, and Xp6 have electron affinities less than zero because they are unusually stable.  In other words instead of energy being given off, these elements actually require an input of energy in order to gain electrons. e.g. Be, N, Ne *Proprietary material of SILVER OAK UNIVERSITY Page | 11  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah e. Electronegativity:  The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is known as electronegativity.  Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge.  Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not form compounds, they do not have electronegativities. Note that there is little variation among the transition metals. Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.  Note that there is little variation among the transition metals.  Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.  Of the main group elements, Fluorine has the highest electronegativity (EN = 4.0) and Caesium the lowest (EN = 0.79).  This indicates that fluorine has a high tendency to gain electrons from other elements with lower electronegativities. Electronegativities can be used to predict what happens when certain elements combine. *Proprietary material of SILVER OAK UNIVERSITY Page | 12  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 2.2.2 Trends in Chemical Properties: a. Periodicity of Valence or Oxidation States:  While moving left to right across a period, the number of valence electrons of elements increases and varies between 1 and 8.  But the valency of elements, when combined with H or O first, increases from 1 to 4 and then it reduces to zero.  Consider two compounds containing oxygen Na2O and F2O. In F2O, the electronegativity of F is more than oxygen. Hence, each of F atoms will attract one electron from oxygen i.e. F will show -1 oxidation state and O will show +2 oxidation state.  Whereas, in the case of Na2O, oxygen is highly electronegative than sodium atom. So oxygen will attract two electrons from each sodium atom showing -2 oxidation state and Na will have +1 oxidation state.  The oxidation state of the element represents the charge possessed by an atom due to the loss or gain of electrons in the molecule.  As we move down in a group the number of valence electrons does not change. Hence, all the elements of one group have the same valency. b. Anomalous Properties of Second Period Elements:  It has been observed that Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, and Fluorine have slightly different periodic properties than the rest of the elements belonging to Group 1, 2, 13-17 respectively.  Lithium and Beryllium form covalent compounds, whereas the rest of the members of Groups 1 and 2 form ionic compounds.  The oxide that is formed by Beryllium when it reacts with Oxygen is amphoteric in nature, unlike other Group 2 elements that form basic oxides.  Yet another example is that of Carbon which can form stable multiple bonds, whereas Si = Si double bonds are not very common.  Second-period elements display periodic properties that are similar to the second element of the next group (i.e. Lithium is similar to Magnesium and Beryllium to Aluminium) or in other words, they have a diagonal relationship. Reasons for Anomalous Periodic Properties:  Small size of these atoms  High electronegativity *Proprietary material of SILVER OAK UNIVERSITY Page | 13  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  Large charge/radius ratio  These elements also have only 4 valence orbitals available (2s and 2p) for bonding as compared to the 9 available (3s, 3p, and 3d) to the other members of the respective groups, so their maximum co-valency is 4. (This is why Boron can only form [BF4]– whereas Aluminium can form [AlF6]3-). 2.3 Chemical Bonding:  Molecules of chemical substances are made of two or more atoms joined together by some force, acting between them. This force which results from the interaction between the various atoms that go to form a stable molecule, is referred to as a Chemical Bond.  5 types of Chemical Bonds: 1. Ionic / Electrovalent Bond 2. Covalent Bond 3. Coordinate Covalent Bond 4. Metallic Bond 5. Hydrogen Bond 2.3.1 Ionic / Electrovalent Bond:  Established by transfer of an electron from one atom to another.  The electrostatic attraction between the cation (+) and anion (–) produced by electron-transfer constitutes an Ionic or Electrovalent bond.  The compounds containing such a bond are referred to as Ionic or Electrovalent Compounds. Conditions for formation of Ionic Bond: *Proprietary material of SILVER OAK UNIVERSITY Page | 14  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 1. Number of valence electrons 2. Net lowering of Energy 3. Electronegativity difference of A and B Examples of Ionic Compounds:  NaCl  CaO  MgCl2  Al2O3 Characteristics of Ionic Compounds:  Solids at Room Temperature  High Melting Points  Hard and brittle  Soluble in water  Conductors of electricity  Do not exhibit isomerism  Ionic reactions are fast 2.3.2 Covalent Bond:  Two atoms could achieve stable 2 or 8 electrons in the outer shell by sharing electrons between them.  The shared pair is indicated by a dash (–) between the two bonded atoms. A shared pair of electrons constitutes a Covalent bond or Electron-pair bond.  The attractive force between atoms created by sharing of an electron-pair.  The compounds containing a covalent bond are called covalent compounds. Conditions for formation of Covalent Bond: *Proprietary material of SILVER OAK UNIVERSITY Page | 15  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 1. Number of valence electrons 2. Equal electronegativity 3. Equal sharing of electrons Examples of Covalent Compounds:  H2  Cl2  H2O  NH3  O2  N2  CO2 Characteristics of Covalent Compounds:  Gases, liquids or solids at room temperature  Low melting points and boiling points  Neither hard nor brittle  Soluble in organic solvents  Non-conductors of electricity  Exhibit Isomerism  Molecular reactions 2.3.3 Coordinate Covalent Bond:  A coordinate covalent bond is formed when both the electrons are supplied entirely by one atom.  The atom A which donates the lone pair is called the donor, while B which accepts it the acceptor. *Proprietary material of SILVER OAK UNIVERSITY Page | 16  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  The molecule or ion that contains the donor atom is called the ligand. Examples of Covalent Compounds:  NH4+  H3O+  BF4-  Al2Cl6 (AlCl3)  O3 2.3.4 Metallic Bond:  Collective sharing of a sea of valence electrons between several positively charged metal ions. Characteristics of Metallic Compounds:  Good conductor of electricity  Good conductor of heat  Highly ductile  Shiny metallic lustre  High Melting and Boiling Points Examples of Metallic Compounds:  Iron, Cobalt, Calcium, Silver, Gold, Barium, Platinum, Chromium, Copper, Zinc, etc. *Proprietary material of SILVER OAK UNIVERSITY Page | 17  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 2.3.5 Hydrogen Bond:  Electrostatic force of attraction between a hydrogen (H) atom which is covalently bound to a more electronegative atom or group, particularly the second-row elements nitrogen (N), oxygen (O), or fluorine (F). *Proprietary material of SILVER OAK UNIVERSITY Page | 18  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  There are two types of H bonds, and it is classified as the following: 1. Intermolecular Hydrogen Bonding  When hydrogen bonding takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding.  For example – hydrogen bonding in water, alcohol, ammonia etc. 2. Intramolecular Hydrogen Bonding  The hydrogen bonding which takes place within a molecule itself is called intramolecular hydrogen bonding.  It takes place in compounds containing two groups such that one group contains hydrogen atom linked to an electronegative atom and the other group contains a highly electronegative atom linked to a lesser electronegative atom of the other group.  For example – hydrogen bonding in o – nitrophenol Significance of Hydrogen Bonding:  Water evaporates slowly from the earth.  Flexibility of muscles  Water retain inside the soil  Solubility of substances  Effectiveness of drugs  Strength of cement concrete  Elasticity of nylon polymer 2.4 Mole Concept: 2.4.1 Mole:  One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) of the 12C isotope. *Proprietary material of SILVER OAK UNIVERSITY Page | 19  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  It is known as ‘Avogadro constant’, denoted by NA in honour of Amedeo Avogadro. 2.4.2 Molar Mass:  The mass of one mole of a substance in grams is called its molar mass.  Molar mass of Water = 18 g mol-1  Molar mass of Sodium Chloride = 58.5 g mol-1  Molar Mass of Hydrochloric Acid = ____________  Molar Mass of Sulfuric Acid = ____________ 2.4.3 Mass percentage of Element:  Mass % of Elements (H & O) in Water = __________  Mass % of Elements (Na & Cl) in Sodium Chloride = __________  Mass % of Elements (H & Cl) in Hydrochloric Acid = __________  Mass % of Elements (H, S & O) in Sulfuric Acid = __________ 2.5 Theory of Solutions: 2.5.1 Types of Solution: Sr. No. State of Solute State of Solvent Example 1 Gas Gas Air Oxygen in Water, Carbonated drinks 2 Gas Liquid (i.e. CO2 in water) 3 Gas Solid Adsorption of H2 by Palladium 4 Liquid Liquid Alcohol in water 5 Liquid Solid Mercury in silver 6 Solid Liquid Sugar + water, Salt + water 7 Solid Solid Metal alloys (i.e. Steel) 2.5.2 Concentration of Solution:  The amount of solute present in a given amount of solution. *Proprietary material of SILVER OAK UNIVERSITY Page | 20  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  A solution containing a relatively low concentration of solute is called dilute solution. A solution of high concentration is called concentrated solution. Ways of expressing Concentration: 1. Per cent by weight (% w/w):  Weight of the solute as a per cent of the total weight of the solution.  For example, if a solution of HCl contains 36 per cent HCl by weight, it has 36 g of HCl for 100 g of solution. 2. Mole Fraction:  The ratio of the number of moles of solute and the total number of moles of solute and solvent.  If n represents moles of solute and N number of moles of solvent,  Mole fraction of solvent would be,  Mole fraction is unitless. *Proprietary material of SILVER OAK UNIVERSITY Page | 21  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 3. Molarity:  The number of moles of solute per litre of solution.  If n is the number of moles of solute and V litres the volume of solution,  Unit: mol litre-1 ( )  Denoted by: M 4. Molality:  The number of moles of solute per kilogram of solvent  Molality is defined in terms of mass of solvent while molarity is defined in terms of volume of solution. 5. Normality:  Number of equivalents of solute per litre of the solution  Denoted by: N Give it a try: *Proprietary material of SILVER OAK UNIVERSITY Page | 22  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 1. What is the per cent by weight of NaCl if 1.75 g of NaCl is dissolved in 5.85 g of water? 2. Calculate the mole fraction of HCl in a solution of hydrochloric acid in water, containing 36 per cent HCl by weight. 3. What is the molarity of a solution prepared by dissolving 75.5 g of pure KOH in 540 ml of solution? 4. What weight of HCl is present in 155 ml of a 0.540 M solution? 5. What is the molality of a solution prepared by dissolving 5.0 g of toluene (C7H8) in 225 g of benzene (C6H6)? 6. 5 g of NaCl is dissolved in 1000 g of water. If the density of the resulting solution is 0.997 g/ml, calculate the Normality. (Assume volume of the solution is equal to that of solvent.) 2.6 Structure of Material (Solids):  There are mainly four type of arrangements are found in solid materials. 2.6.1 Ionic Solids:  Ionic solids are made up of positive and negative ions.  In this type of solids, positive and negative ions are surrounded by each other.  Examples: Structures of NaCl (Figure 1 – LHS), CsCl (Figure 2 – RHS), etc. 2.6.1 Metallic Solids:  Metallic solids contain metallic bonds in their molecular structure.  Metallic solids contain FCC, BCC or HCP arrangement. *Proprietary material of SILVER OAK UNIVERSITY Page | 23  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  Examples: FCC = Al, Ni, Ag, Cu, etc., BCC = Fe, V, Nb, Cr, etc. HCP = Ti, Zn, Mg, Cd, etc. 2.6.3 Molecular Solids:  In case of molecular solids, atoms in molecule are joined by covalent bonds.  These types of solid materials consist of Van Der Waal’s force of attraction between constituent covalent molecules of molecular solids.  Examples: P4, S8, etc. 2.6.4 Network Solids:  In these solids, covalent bond is distributed throughout lattice.  Examples: Diamond, Graphite, Asbestos, Mica, Quartz, etc.  Diamond and Graphite are the allotropes of Carbon. a. Diamond: *Proprietary material of SILVER OAK UNIVERSITY Page | 24  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  In diamond, each carbon is bonded to four other carbon atoms and form sp3 hybridization.  It has tetrahedral structure.  It requires more energy to break the bond; hence its melting point is very high (i.e. around 3500 oC). b. Graphite:  In graphite, each carbon atom forms bonds with three other carbon atoms in plane and forms sp2 hybridization.  It has continuous hexagonal network arrangement. c. Asbestos:  One Si atom joins with two oxygen atoms and with two other mono valent atoms forming a stable long chain. d. Mica: *Proprietary material of SILVER OAK UNIVERSITY Page | 25  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah  One Si atom combines with three oxygen atoms and with one other mono valent atom and forms stable molecule of Mica.  It is a two dimensional network solid having planner shape. e. Quartz:  One Si atom combines with four oxygen atoms forming stable molecule of Quartz.  Quartz is a three dimensional network solid having tetrahedral shape.  It has very high melting point. 2.7 Crystallization of Inorganic Salts: 1. The solution is heated in an open container. 2. The solvent molecules start evaporating, leaving behind the solutes. 3. When the solution cools, crystals of solute start accumulating on the surface of the solution. 4. Crystals are collected and dried as per the product requirement. 5. The undissolved solids in the liquid are separated by the process of filtration. 6. The size of crystals formed during this process depends on the cooling rate. 7. Many tiny crystals are formed if the solution is cooled at a fast rate. 8. Large crystals are formed at slow cooling rates. *Proprietary material of SILVER OAK UNIVERSITY Page | 26  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah Activity:  Take 50 ml water in a beaker  Add sugar in it and stir it  Now heat the solution  Repeat the process continuously  After some time there will be a point at which no more sugar can be dissolved in water. This stage is the saturation point, and the solution is referred to as a saturated solution  Now filter the sugar with the help of a filter paper  Collect the filtrate in a glass bowl and cool it  We will observe that some fine crystals are formed in the bowl  The process of filtration can separate these crystals. The liquid left after the removal of crystals is known as mother liquor Check Your Knowledge: 1. Ionic Radius of Be2+ is ___________ compared to that of the O2-. (A) More (C) Same (B) Less 2. As we move across a period, electron affinity, __________________ (A) Decreases (C) Increases (B) Remains constant (D) Can’t say anything 3. Which of the following has the lowest value of EN? (A) Caesium (C) Fluorine (B) Selenium (D) Lithium 4. ____________ type of bond is formed when both the electrons are supplied entirely by one atom. (A) Covalent (C) Ionic (B) Coordinate Covalent (D) Hydrogen *Proprietary material of SILVER OAK UNIVERSITY Page | 27  SUBJECT (CHEMISTRY) LECTURE COMPANION SEMESTER: I/II PREPARED BY: Mr. Nishit Shah 5. Which of the following is not from the same category according to the arrangement of the atoms? (A) Zinc (C) Silver (B) Gold (D) Copper Answers: (1) B (2) C (3) A (4) B (5) A -------- X ------------------ X ------------------ X ------------------ X ------------------ X -------- *Proprietary material of SILVER OAK UNIVERSITY Page | 28