Grade 11 Chemistry Past Paper 2016 Unit 2 - PDF

 UNIT 2 Chemical Bonding and Structure Chemical bond is: A force that holds two or more atoms together. A lasting attraction between atoms or ions that enables the formation of molecules and crystals. Octet rule( rule of eight) states that:...

 UNIT 2 Chemical Bonding and Structure Chemical bond is: A force that holds two or more atoms together. A lasting attraction between atoms or ions that enables the formation of molecules and crystals. Octet rule( rule of eight) states that: during the formation of a chemical compound, each atom has octet (8) electrons in its highest occupied energy level by gaining, losing, or sharing electrons. A tendency of atoms to gain or lose electrons so as to form ions that are isoelectronic with the nobel gases Intramolecular forces are the forces that hold atoms together within a molecule. Intermolecular forces are forces that exist between molecules. Limitation of octet rule: Rule of two is one of the exceptions of octet rule The ions of some transition elements do not have the usual noble gas valence shell of ns2np6 and are not isoelectronic with any of the noble gases The ions of the post transition do not usually obey the octet rule. NOTE! Post transition elements are allocated near to the right of the transition group Types of Chemical Bonding There are three main types of chemical bonds : Covalent ( electron pair bonding ) Ionic bonding ( electrovalent bonding ) metallic bonding Ionic Bonding is formed by electron transfer from a metal to a non-metal with different electronegativity values. Note! Ionic compounds are usually formed when metal cations bond with non-metal anions. The only common exception is ammonium ion which is not a metal, but it forms ionic compounds. Compounds in which particles are held together by ionic bonds are called ionic compounds. Covalent Bonding is formed as a result of electron sharing between two non-metals. If the electronegativity values are very similar then it is non-polar covalent bonding but if the electronegativity values are much different, then it is a polar covalent bonding. Metallic Bonding refers to the interaction between the delocalized electrons and the metal nuclei. Lewis Electron-Dot Symbols: It is a representation of an atom or monoatomic ion showing the valence electrons, if present, as dots placed around the letter symbol of the element. Formation of Ionic Bonding:  The formation of ionic compounds is not merely the result of low ionization energies and high affinities for electrons, although these factors are very important.  It is always an exothermic process; the compound is formed because it is more stable (lower in energy) than its elements.  Much of the stability of ionic compounds result from the packing of the oppositely charged positive and negative ions together.  A measure of just how much stabilization results from this packing is given by the lattice energy (U).  The lattice energy is an important indication of the strength of ionic interactions and is a major factor influencing melting points, hardness, and solubility of ionic compounds. NOTE: Formation of ionic bonding is affected by  Ionization energy  Electron affinity  Lattice energy NB:  Lattice energy is the energy released when one mole of cat ion and one mole of an ions  Combine together to form one mole of ionic compounds.  It is also the energy required to break one mole of ionic solid in to isolated gaseous ions  It affects strength, melting point, hardness and solubility of ionic compounds  The higher the lattice energy the higher is the bond strength, melting point and hardness but the lower the solubility of ionic compounds. NOTE:  Lattice energy is affected by size ion and the magnitude of the charge.  Lattice energy is a measure of the strength of the ionic bonds in an ionic compound.  It provides insight into several properties of ionic solids including their volatility, their solubility, and their hardness.  The lattice energy of an ionic solid cannot be measured directly.  However, it can be estimated with the help of the Born-Haber cycle.  Generally, this quantity is expressed in terms of kilojoules per mole (kJ/mol). Comparison between Lattice Energy and Lattice Enthalpy The molar lattice energy of an ionic crystal can be expressed in terms of molar lattice enthalpy, pressure, and change in volume via the following equation: ΔLatticeU = ΔLatticeH – pΔVm Where:  ΔLatticeU denotes the molar lattice energy.  ΔLatticeH denotes the molar lattice enthalpy.  ΔVm is the change in volume (per mole).  p is the pressure. Lattice energy is calculated by using Hess’s law through the born Haber cycle. Born Haber cycle is a cycle of enthalpy change of process that leads to the formation of a solid crystalline ionic compound from the elemental atoms in their standard state and of the enthalpy of formation of the solid compound such that the net enthalpy becomes zero. Hess’s law: states that the total enthalpy change for a reaction equals the sum of the enthalpy changes for all intermediates steps in the reaction. Factors Affecting Formation of Ionic Bonding  Ionization energy (IE): Elements having low IE have a more favorable chance to form a cations, thereby having a greater tendency to form ionic bonds.  Electron affinity (EA): The elements having higher electron affinity favour the formation of an ionic bond.  Lattice energy: Larger lattice energy would favor the formation of an ionic bond. Exceptions to Octet Rule in Ionic Compounds:  Less than Octet (Central Atom is Deficient of Electrons)  More than Octet (18-Electron Rule) Properties of Ionic Compounds  They are usually crystalline solids.  They have high melting points and high boiling points.  They are usually soluble in water but insoluble in organic solvents.  They conduct electricity when dissolved in water or when melted.  They are aggregates of ions and non-volatile.  They are very resistant to heat but many will be easily broken by water.  Covalent Bonds and molecular geometry A covalent bond is a chemical bond that involves the sharing of electrons to form electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs. An optimum distance between nuclei, called the bond length. A bonding pair consists of two electrons shared between atoms, creating a bond. A lone pair of an atom consists of two electrons not involved in a bond. Example:  A Lewis Structure is a very simplified representation of the valence shell electrons in a molecule.  It is used to show how the electrons are arranged around individual atoms in a molecule. Electrons are shown as "dots" or for bonding electrons as a line between the two atoms  Coordinate Covalent Bonds A covalent bond in which both electrons are donated by one of the atoms is called a coordinate covalent bond (dative bond) or donor acceptor bond Example:  Resonance structure Resonance structures are sets of Lewis structures that describe the delocalization of electrons in a polyatomic ion or a molecule. The resonance hybrid is more stable than any individual resonance structures. Features of resonance structure:  All resonance structures must have the same atom connectivity and only differ in the electron arrangement.  (Atoms NEVER move; only electrons move.)  All resonance structures have the same number of electrons and net charge. Example:  Exceptions to the Octet rule in Covalent Bonding many Lewis structures follow the octet rule, there are exceptions. There can be categorized into three groups: These are  Less than octet (central atom is deficient of electrons):  More than octet (central atom has excess of electrons):  Molecules containing an odd number of electrons  Polar and Non-Polar Covalent Molecules The main difference between polar and nonpolar covalent bonds is that polar covalent bonds are chemical bonds where electrons are shared unequally between two atoms due to the difference in electronegativity, whereas nonpolar covalent bonds are chemical bonds in which electrons are shared equally between two atoms due to the similar electro negativities.  A dipole moment is the product of the magnitude of the charge and the distance between the centers of the positive and negative charges. It is denoted by the Greek letter 'µ'.  The dipole moment is a vector measure whose direction runs from negative to a positive charge.  The formula for electric dipole moment for a pair of equal & opposite charges is p = qd, the magnitude of the charges multiplied by the distance between the two.  A molecule that carries opposite charges is called dipole.  If dipole moment is zero the molecule is non polar  For polar molecule the dipole moment is not equal to zero.  The net dipole moment of a molecule arises as the sum of individual bond dipoles.  When the electronegativity difference between two atoms become greater, the bond becomes more polar.  The SI unit of dipole moment is coulomb meter  Other unit of dipole moment is Debye. 1D = 3.33×10-30 C.m Example: HCl - polar CH4 - non polar  Properties of Covalent Compounds Covalent compounds share some common properties:  Low melting points.  Low boiling points.  Poor electrical conductors.  Poor thermal conductors.  Form brittle or soft solids.  Low enthalpies of fusion.  Low enthalpies of vaporization.  Covalent compounds exist as separate molecules Due to weak intermolecular forces, many covalent molecules or covalent  compounds are liquids or gases at room temperature. However, some covalent  molecules like iodine are solids at room temperature.  Covalent compounds are volatile.  Covalent compounds are generally insoluble in water. Most covalent compounds  are soluble in non-polar solvents.  Non-polar covalent compounds are non-electrolytes because they do not  conduct electricity.  Molecular Geometry The spatial arrangement of atoms in a molecule is called molecular geometry or molecular shape. The molecular part is determined by the mutual repulsion of valence electrons about an atom in a molecule. It is determined by the central atom and the surrounding atoms and electron pairs. The shape of most molecules can be predicted using the Valence Shell Electron Pair Repulsion (VSEPR) method. Based on the VSEPR model we can predict:  The arrangement of electron pairs about the central atom  The molecular geometry In general, according to the VSEPR model, the repulsive forces decrease in the following order: Lone pair vs lone pair > lone pair vs bonding pair > bonding pair vs bonding pair The 5 molecular geometries are linear, trigonal planar, tetrahedral, trigonal pyramidal and octahedral -To keep track of the total number of bonding pairs and lone pairs, we show molecules with lone pairs as ABx E y , where A is the central atom B is a surrounding atom E is a lone pair on Electron pair geometry or electron pair arrangement is the set of electrons around the central atom A set of electron means an electron domain One set of electron may be one single bond, one double bond, one triple bond, one lone pair or one electron Types of electron pair arrangement  Linear - electron pair arrangement Two set of electrons around the central atom  Trigonal planar - three set of electrons around the central atom  Tetrahedral - four set of electrons around the central atom  Trigonal bi pyramidal - five set of electrons around the central atom  Octahedral – six set of electrons around the central atom  Molecular Shape and Molecular Polarity The overall polarity of molecules with more than one bond is determined from both the polarity of the individual bonds and the shape of the molecule. Each bond's dipole moment can be treated as a vector quantity, having a magnitude and direction. Bond Polarity and Dipole Moment  A bond dipole moment is a measure of the polarity of a chemical bond between two atoms in a molecule.  It involves the concept of electric dipole moment, which is a measure of the separation of negative and positive charges in a system  Bond polarity refers to the distribution of electric charge across a chemical bond between two atoms.  If the bond is non-polar, the charge is evenly distributed across the bond.  If the bond is polar, one end of the bond will have a slightly positive charge and the other end will have a slightly negative charge.  Intermolecular force  Intermolecular forces, often abbreviated to IMF, are the attractive and repulsive forces that arise between the molecules of a substance.  These forces mediate the interactions between individual molecules of a substances  Intermolecular forces are responsible for most of the physical and chemical properties of matter  The three types of attractive force are known to exist between neutral molecules: dipole– dipole forces, London (or dispersion) forces, and hydrogen bonding forces.  The term van der Waals forces are a general term for those intermolecular forces that include dipole–dipole and London forces.  Van der Waals forces are the weak attractive forces in a large number of substances Dipole-Dipole forces  Dipole-dipole forces act between the molecules possessing permanent dipole.  When polar molecules are brought near one another, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces; the partially positive end of one molecule attracts the partially negative end of another.  Ends of the dipoles possess “partial charges” and these charges are shown by Greek letter delta (δ).  Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Hydrogen bonding  Hydrogen bonding is a special type of dipole-dipole attraction between molecules, not a covalent bond to a hydrogen atom.  It results from the attractive force between a hydrogen atom covalently bonded to a very electronegative atom such as a N, O, or F atom and another very electronegative atom. Dispersion force  The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles  The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles  Metallic Bonding Metallic bonds are the chemical bonds that hold atoms together in solid metals. In these metals, each metal atom is bonded to several neighboring atoms. Metallic bonding is the sharing of free electrons (delocalized electrons) among a lattice of positively charged metal ions. The bonding electrons are relatively free to move throughout the three-dimensional structure. The strength of the metallic bond depends on: 1. the number of electrons in the delocalized “sea” of electrons. More delocalized electrons result in a stronger bond and a higher melting point. 2. the packing arrangement of the metal atoms. The more closely packed the atoms are the stronger the bond is and the higher the melting point.  Electron sea model  The electron-sea model is a very simple model, which pictures the metal as an array of metal cations in a “sea” of electrons  It proposes that all the metal atoms in a sample pool their valence electrons to form an electron “sea” that is delocalized throughout the piece.  The metal ions (nuclei plus core electrons) are submerged within this electron sea in an orderly array.  This model proposes that all the metal atoms in a metallic solid contribute their valence electrons to form a "sea" of electron.  The electrons present in the outer energy levels of the bonding metallic atoms are not held by any specific and can move easily from one atom to the next.  Properties of metals and Bonding  The general properties of metals include malleability , ductility , most are strong and durable, good conductors of heat and electricity, atoms are difficult to separate, but malleability and ductility suggest that the atoms are relatively easy to move in various directions.  Metallic bonds show typical metallic properties such as high electrical conductivity, lustre, and high heat conductivity.  Metals are good conductors of electricity and heat because of their mobile electrons.  They are strong and opaque in nature.  Most metals are solids with high melting and much higher boiling points; because the atoms of metals have strong attractive forces between them and much energy is required to overcome this force. The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding.  Chemical Bonding Theories There are 2 bonding theories in chemistry: Valence bond theory: chemical bonds are formed when atomic orbitals overlap. Molecular orbital theory: quantum mechanical treatment of bonding describing the electronic structure of molecules. Valence Bond Theory (VBT)  Lewis theory provides a simple, qualitative way to describe covalent bonding.  VSEPR theory allows predicting the probable molecular shapes.  The basic principle of valence bond theory is that a covalent bond forms when orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.  The overlap results in an increased electron charge-density in the region between the atomic nuclei.  The increased density of negative charge serves to hold the positively-charged atomic nuclei together.  In the valence bond (VB) method, a covalent bond is a region of high electron charge- density that results from the overlap of atomic orbitals between two atoms.  The strength of the bond between any two atoms generally depends on the extent of the overlap between the two orbitals.  As the two atoms are brought more closely together, however, the repulsion of the atomic nuclei becomes more important than the electron-nucleus attraction and the bond become unstable. As the two atoms are brought more closely together, however, the repulsion of the atomic nuclei becomes more important than the electron-nucleus attraction and the bond become unstable. For each bond, then, there is a condition of optimal orbital overlap that leads to maximum bond strength (bond energy) at a particular internuclear distance (bond length). Note! The s orbital is spherical, but p and d orbitals have particular orientations - more electron density in one direction than in another - so a bond involving p or d orbitals will tend to be oriented in the direction that maximizes overlap. Various type of atomic orbital overlap leads to covalent bond formation. Three simple basic ones are  s-s overlaps in which half-filled s orbitals overlap,  s-p overlaps where half-filled s orbital of one atom overlaps with one of the p orbital having one electron only  p-p overlap in which two half-filled p orbitals overlap.  Hybridization of Orbitals  Valence bond theory employs the concept of hybridization. It is the overlap or blending of s, p and d orbitals to explain bond formation.  Hybridization is the idea that atomic orbitals fuse to form newly hybridized orbitals, which in turn, influences molecular geometry and bonding properties.  Hybridization is also an expansion of the valence bond theory.  Hybridization is defined as the process of combining two atomic orbitals to create a new type of hybridized orbitals.  The intermixing of atomic orbitals typically results in the formation of hybrid orbitals with completely different energies, shapes, and so on.  Hybridization is primarily carried out by atomic orbitals of the same energy level. However, both fully filled and half-filled orbitals can participate in this process if their energies are equal.  The concept of hybridization is an extension of valence bond theory that helps us understand bond formation, bond energies, and bond lengths.  Hybridization scheme, the number of new hybrid orbitals is equal to the total number of atomic orbitals that are combined.  Hybrid orbitals are the new orbitals formed as a result of hybridization  Steps to determine the type of Hybridization To understand the type of hybridization in an atom or an ion, the following rules must be followed. Write a reasonable Lewis structure for the species Use VSEPR theory to predict the electron-set arrangement of the central atom Select the hybridization scheme that correspond to the VSEPR prediction Describe the orbital overlap Features of Hybridization Hybridization occurs between atomic orbitals with equal energies. The number of hybrid orbitals formed equals the number of atomic orbitals that mix. It is not required for all half-filled orbitals to participate in hybridization. Even orbitals that are completely filled but have slightly varying energy can participate. Hybridization occurs only during bond formation, not in a single gaseous atom. If the hybridization of the molecule is known, the molecule’s shape can be predicted. The larger lobe of the hybrid orbital is always positive, while the smaller lobe on the opposite side is always negative. Types of Hybridization sp Hybridization  It occurs when one s and one p orbital in an atom’s main shell combine to form two new equivalent orbitals.  The newly formed orbitals are known as sp hybridized orbitals.  It produces linear molecules at a 180° angle.  It entails combining one’s orbital and one ‘p’ orbital of equal energy to produce a new hybrid orbital known as an sp hybridized orbital.  It’s also known as diagonal hybridization.  Each sp hybridized orbital contains the same amount of s and p characters. sp2 Hybridization  It occurs when one s and two p orbitals of the same atom’s shell combine to form three equivalent orbitals.  The newly formed orbitals are known as sp2 hybrid orbitals.  It’s also known as trigonal hybridization.  It entails combining one’s orbital with two ‘p’ orbitals of equal energy to create a new hybrid orbital known as sp2.  A trigonal symmetry mixture of s and p orbitals is kept at 120 degrees.  All three hybrid orbitals remain in the same plane and form a 120° angle with one another.  Each hybrid orbital formed has a 33.33 % and a 66.66 % ‘p’ character.  The molecules with a triangular planar shape have a central atom that is linked to three oter atoms and is sp2 hybridized. sp3 Hybridization  It occurs when one ‘s’ orbital and three ‘p’ orbitals from the same shell of an atom combine to form four new equivalent orbitals.  The hybridization is known as tetrahedral hybridization or sp3.  The newly formed orbitals are known as sp3 hybrid orbitals.  These are pointed at the four corners of a regular tetrahedron and form a 109°28′ angle with one another.  The sp3 hybrid orbitals form a 109.28-degree angle.  Each hybrid orbital has a 25% s character and a 75% p character. sp3d Hybridization  The mixing of 1s orbitals, 3p orbitals, and 1d orbitals results in 5 sp3d hybridized orbitals of equal energy.  Their geometry is trigonal bipyramidal.  The combination of s, p, and d orbitals results in trigonal bipyramidal symmetry. The equatorial orbitals are three hybrid orbitals that are oriented at a 120° angle to each other and lie in the horizontal plane.  The remaining two orbitals, known as axial orbitals, are in the vertical plane at 90 degrees plane of the equatorial orbitals. sp3d2 Hybridization When 1s, 3p, and 2d orbitals combine to form 6 identical sp3d2 hybrid orbitals, the hybridization is called sp3d2 Hybridization. These seven orbitals point to the corners of an octahedron. They are inclined at a 90-degree angle to one another. sp3d3 Hybridization It has 1s, 3p, and 3d orbitals, which combine to form 7 identical sp3d3 hybrid orbitals. These seven orbitals point to the corners of a pentagonal bipyramidal. e.g. IF6. Shapes of Hybridization  Linear: The sp hybridization is caused by the interaction of two-electron groups; the orbital angle is 180°.  Trigonal planar: Three electron groups are involved, resulting in sp2 hybridization; the orbitals are 120° apart.  Tetrahedral: Four electron groups are involved, resulting in sp3 hybridization; the orbital angle is 109.5°.  Trigonal bipyramidal: Five electron groups are involved, resulting in sp3d hybridization; the orbital angles are 90° and 120°.  Octahedral: Six electron groups are involved, resulting in sp3d2 hybridization; the orbitals are 90° apart.  Molecular Orbital Theory (MOT)  Molecular orbital theory is a method of accounting for covalent bonds that depends on quantum theory and mathematical principles.  This theory is based on the fact that electrons are not the substantive little dots as we portray in Lewis structures.  Atomic orbitals are capable of combining or overlapping, to produce new electron distributions called molecular orbitals (MOs) – one MO for every AO.  The quantum mechanical treatment of electrons in atoms as matter waves yields atomic orbitals (AOs).  A similar treatment applied to electrons in molecules yields molecular orbitals (MOs), which are mathematical descriptions of regions of high electron charge density in a molecule. Bonding and Anti-Bonding Molecular Orbitals  Bonding molecular orbitals have a region of high electron density between the nuclei.  Anti-bonding molecular orbitals have region of zero electron density (a node) between the nuclei.  In place of atomic orbitals of the separated atoms, the molecular orbitals for the united atoms are obtained, and these are of two types.  One type, a bonding molecular orbital, places a high electron charge density in between the two nuclei and the other type, an anti-bonding molecular orbital, places a high electron charge density away from the region between the two nuclei. Bond Order The term bond order is used to indicate whether a covalent bond is single (bond order = 1), double (bond order = 2) or triple (bond order = 3). Bond order = 0.5 [(number of electrons in bonding MOS ) - (number of electrons in antibonding MOS)] Magnetic Properties  A species with unpaired electrons exhibits paramagnetic property.  The species is attracted by an external magnetic field.  A species in which all the electrons are paired, exhibits diamagnetism.  Such species are not attracted (and, in fact, are slightly repelled) by a magnetic field. TYPES OF CRYSTALS  A crystal is a piece of a solid substance that has plane surface, sharp edges, and a regular geometric shape.  The fundamental units-atoms, ions or molecules are assembled in a regular, repeating manner extending in three dimensions throughout the crystal.  An essential feature of a crystal is that its entire structure can be figured from a tiny portion of it.  Some solids, like glass, lack this long-range order and are said to be amorphous.  The structural units of an amorphous solid, whether they are atoms, molecules or ions occur at random positions.  In liquids, there is no ordered pattern to the arrangement of an amorphous solid.  A structural unit of a crystalline solid has a characteristic repetitive pattern. There are four important classes of crystalline solids. 1. Ionic Crystals  The fundamental units of an ionic solids are positive and negative ions.  As a result, the inter-particle forces (ionic bonds) are much stronger than the van der Waals forces in molecular solids.  To maximize attractions, cations are surrounded by as many anions as possible, and vice versa.  The properties of ionic solids are direct consequences of the strong inter-ionic forces, which create a high lattice energy.  Crystalline ionic solids are usually, brittle, and non- conductors of electricity, although molten crystals may be good conductors.  They usually have high melting points. 2. Molecular Crystals  Various combinations of dipole-dipole, dispersion and hydrogen-bonding forces are operative in molecular solids, which accounts for their wide range of physical properties. Dispersion forces are the principal forces acting in non-polar substances, so melting points generally increase with molar mass. Among polar molecules, dipole dipole forces and where ever possible, hydrogen-bonding dominate.  Nevertheless, intermolecular forces are still relatively weak, so the melting points are much lower than ionic, metallic and network covalent solids.  The fundamental unit of a molecular solid is the molecule. Such solids are common among organic compounds and simple inorganic compounds.  Molecular crystals are usually transparent, brittle, and break easily when stressed.  They are usually non-conductors of heat and electricity and usually have low melting points. 3. Covalent Network Crystals  In this type of crystalline solids, separate particles are not present.  Strong covalent bonds link the atoms together throughout the network of covalent solid.  As a consequence of the strong bonding, all these substances have extremely high melting  and boiling points, but their conductivity and hardness depend on the nature of their bonding.  These crystals are usually hard, non-conductors of heat and electricity, and have high melting points.  4. Metallic Crystals  The strong metallic bonding forces hold individual atoms together in metallic solids.  The fundamental units of pure metallic solids are identical metal atoms.  Metallic crystals are opaque with reflective surfaces.  They are ductile and malleable, good conductor of heat and electricity, and they usually have high melting points.