Structure and Properties of Ionic and Covalent Substances PDF

Structure and properties of ionic covalent and metallic substances 1.CRYSTAL LATTICE STRUCTURE OF IONIC SOLIDS 2. ALLOTROPES 3. SIMPLE MOLECULAR SOLIDS 4. GIANT COVALENT SOLIDS( GIANT MOLECULAR CRYSTALS) Objectives:  Describe structure of ionic crystals and give examples.  Relate the pro...

Structure and properties of ionic covalent and metallic substances 1.CRYSTAL LATTICE STRUCTURE OF IONIC SOLIDS 2. ALLOTROPES 3. SIMPLE MOLECULAR SOLIDS 4. GIANT COVALENT SOLIDS( GIANT MOLECULAR CRYSTALS) Objectives:  Describe structure of ionic crystals and give examples.  Relate the properties and use of sodium chloride to its structure.  Make diagrammatic representations of sodium chloride.  Make models of sodium chloride in groups for grading.  Describe structure of simple molecular crystals and give examples of each  Distinguish between ionic and simple molecular crystals (solids) by use of melting point, solubility in water and organic solvents, and thermal and electrical conductivity.  Investigate melting point, solubility of solids and conductivity of resulting solutions for lab 1.  Describes structure of giant molecular crystals and give examples.  Explain the term allotropy with reference to the allotropes of carbon – diamond and graphite.  Make diagrammatic representations of graphite and diamond.  Relate the properties and uses of diamond and graphite to their structure. Ionic, simple molecular and giant molecular crystals  A crystal is a solid substance, the structure of which has an ordered arrangement of atoms, molecules or ions. This is called a crystal Lattice structure.  We can classify the structure of such solids as either (1) ionic crystal, (2) Simple molecular crystal and ( 3) Giant molecular crystal. The distinction is dependent on the type of bonding present( whether ionic or covalent as well as the arrangement of the atoms or molecules in the compounds.  Solids that do not have a crystalline structure are called amorphous, e.g. rubber and plastic. Ionic crystals/ Ionic substanc es The Ionic lattice is also called Crystal lattice Structure and properties ionic substances eg. of sodium chloride Uses of NaCl also include use as an electrolyte(An electrolyte is a substance that produces an electrically conducting solution when dissolved in a polar solvent, such as water).  The properties and uses of sodium chloride can be explained by its structure Strong ionic bonds Crystalline solid polar  Dissolution and solvation of NaCl When dissolution happens, the solute separates into ions or molecules, and each ion or molecule is surrounded by molecules of solvent.  You must learn how to draw the crystalline structure of the ionic solid NaCl for CXC exams The given figure clearly illustrates the unit cell structure of NaCl with 14 Cl– and 13 Na+ ions. Thus, there are four NaCl molecules per unit cell. Each Na+ is surrounded by 6 Cl- ions and each Cl- is surrounded by 6 Na+ ions in repeating fashion. A lattice is an ordered array of points in 2 or 3 dimensions describing the arrangement of particles that form a crystal. The unit cell of a crystal is defined by the lattice points Create a model of the sodium chloride structure: Valued: 50 marks Rubric for grading:  10marks creativity  10marks Accuracy  10marks use of color  10marks: correct amount of Na+ ions and Cl- ions  10 marks for Correct positions of Na+ ions and Cl- ions Please note that you can use matches sticks or icicle pops to represent the bonds. You can use flour dough that is colored with food coloring or play dough to represent the sodium cations and chlorine anions Take pictures and send to me from different angles Due two week time. Friday the 17th Intramolecular forces of attraction vs intermolecular forces of attraction of molecules Weak intermolecular forces Intermolecular forces arise because of the attraction between the dipoles in neighboring molecules. There are three types of intermolecular force: Types of weak intermolecular forces  hydrogen bonding  permanent dipole–dipole  van der Waals.  Strength differences  Intermolecular forces are weak compared with covalent, ionic and metallic bonding. The comparative strength of these intermolecular forces are generally in the order from strongest to weakest:  hydrogen bonding( strongest) > permanent dipole–dipole > van der Waals( weakest). A fine example of hydrogen bonding- the force keeping molecules in a drop of water attracted to each other Permanent dipole–dipole forces  Permanent dipole–dipole forces are the weak attractive forces between the δ+ of the dipole of one molecule and the δ- of the dipole of a neighboring molecule. These form between polar molecules Examples  Hydrogen chloride (HCl), hydrogen fluoride (HF), and water are all examples of dipole-dipole forces (H2O) which are given below: HCl (hydrogen chloride): HCl is a permanent dipole.  van der Waals forces van der Waals forces of attraction are not permanent. All atoms and molecules, including noble gas atoms, have van der Waals forces. 1.Electrons in atoms are always in motion. 2.So the electron density may be greater in one part of the molecule than another. 3.So an instantaneous dipole is formed. 4.This dipole can induce the formation of a dipole in a neighbouring molecule. 5.The two neighbouring molecules attract each other because of their dipoles.  London dispersion force -A type of Vander Waals 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. This force is sometimes called an induced dipole-induced dipole attraction. It is the weakest of all intermolecular forces. Found in many diatomic non-polar molecules( lack dipole). Useful info  The peculiar properties of water, e.g. lower density of ice than water and high surface tension, are due to more extensive hydrogen bonding than in many other molecules. Water molecules have two hydrogen atoms and two lone pairs of electrons. So water can form an average of two hydrogen bonds per molecule. Hydrogen bonding  Hydrogen bonding is a special form of permanent dipole bonding. It requires:  1.one molecule with an H atom covalently bonded to an F, O or N atom. These are the most electronegative atoms  2. a second molecule having a F, O or N atom with a lone pair of electrons Hydrogen bonding in water affects its boiling point. Do you think it would be a high or low boiling point?  Simple molecular crystals Covalently bonded substances fall into two main types:  simple molecular crystals ( Also called simple molecular solids)  giant covalent solids( also called Giant atomic crystals. ) Simple molecules (Simple molecular crystals)contain only a few atoms held together by covalent bonds. An example is carbon dioxide (CO2), the molecules of which contain one atom of carbon bonded with two atoms of oxygen. However, although the covalent bonds holding the atoms together in a simple molecule are strong, the intermolecular forces between the simple molecules are weak. Therefore they have lower boiling point and melting points than giant molecular solids. For e.g HCl melting point is -115 degrees Celsius but for NaCl the melting point is 801 degrees Celsius Example of simple molecular crystals are:  Water (ice).  Solid carbon dioxide( Dry ice).  The element Sulphur( S8 molecule)  The element Phosphorus( P4)- white phosphorus  The element Iodine( I2)  Sucrose(table sugar).  Short chain hydrocarbons such as CH4( methane)  Water (ice).  Solid carbon dioxide CO2( Dry ice). Structure of dry ice solid CO2  Structure of the element Sulphur( S8 molecule)  Sulfur is a simple molecular substances with only weak intermolecular forces of attraction (van der Waals forces of attractions) between the molecules. Hence only little energy is needed to overcome them. Its melting point is 115.2 °C. Boiling point is 444.6 °C. Structure of white phosphorus ( P4) The most important elemental form of phosphorus in terms of applications is white phosphorus. It consists of tetrahedral P4 molecules, in which each atom is bound to the other three atoms by a single bond.’ Phosphorus contains P4 molecules. To melt phosphorus you don't have to break any covalent bonds - just the much weaker van der Waals forces between the molecules. Because of the relatively weak intermolecular attractions (van der Waals forces) between the separate P4 molecules, the solid melts easily at 44.1 °C (111.4 °F) and boils at about 280 °C  Structure of I2 molecule Iodine is a dark grey crystalline solid with a purple vapor. Melting point : 114°C. Boiling Point: 184°C. It is very, very slightly soluble in water, but dissolves freely in organic solvents. Iodine is therefore a low melting point solid. The crystallinity suggests a regular packing of the molecules. The structure is described as face centered cubic - it is a cube of iodine molecules with another molecule at the center of each face. The sublimation of iodine is evidence for the presence of weak intermolecular forces in such solids Question  Why does simple molecular crystals/ simple molecular substances have low melting and boiling point?  What are some examples of simple molecular cystals? Differentiating between the structure of ionic and covalent substances  Properties of simple molecular solids No free electrons or ions Summary: Ionic vs Simple covalent ( molecular) substances properties Giant atomic crystals You must be able to draw the structure of diamond and state its properties. You must also be able to use its structure to explain its properties and uses in CXC exam Properties and uses of graphite is left fre You must be able to draw the structure of graphite and state its properties. You must also be able to use its structure to explain its properties and uses in CXC exam Negatively charged electrode cations migrate to it Questions continued  Allotropy Allotropy or allotrophism is the property of some chemical elements to exist in two or more different forms, in the same physical state, known as allotropes of the elements. Examples of Allotrophes  Allotropes are different structural modifications of an element; the atoms of the element are bonded together in a different manner  The allotropes of carbon are diamond and graphite and Buckminister fullerenes  The allotropes of Sulphur are Rhombic Sulphur( alpha Sulphur) and monoclinic Sulphur ( beta Sulphur)  The allotropes of phosphorus are white phosphorus and red phosphorus  Allotropes of carbon Buckminsterfullerene( Bucky ball) is a type of fullerene with the formula C 60. It has a cage-like fused-ring structure (truncated icosahedron) that resembles a soccer ball made of twenty hexagons and twelve pentagons. Each carbon atom has three bonds. It is a black solid that dissolves in hydrocarbon solvents to produce a violet solution. The compound has received intense study, although few real world applications have been found.  Octahedral shaped needle shaped crystal Allotropes of sulphur  Allotropes of Sulphur are rhombic Sulphur and monoclinic Sulphur Allotropes of phosphorus Black phosphorus has an orthorhombic pleated honeycomb structure and is the least reactive allotrope Red phosphorus exists in a polymeric chain of tetrahedrally structured P4 molecules in which one of the P-P bonds are broken to enable the linking of these tetrahedrons Giant Metallic crystal lattices and characteristics of metals Can you believe that this solid piece of aluminum actually has a lattice arrangement? Lattice structures of metals Metals are arranged in very compact and orderly patterns called a lattice.  Structure of Metals  The metal ions are in a fixed position of a metal lattice and the electrons are free to move throughout the lattice. A metal lattice describes the periodic arrays of a solid.  The bond results from the attraction between the free moving electrons and the positive ions. This force is called the electrostatic force of  Structure of Metals Electrostatic forces of attraction between the positively charged cations and the negatively charged electrons hold the lattice together. A metal is therefore a seen as a rigid framework of cations immersed in a ‘sea’ of electrons that serve as the cement holding the three-dimensional cationic network together – Metallic bonding. Characteristics of metals  Solids at room temperature except for mercury  Malleable  Ductile  High melting and Boiling points  High density  Good conductors of heat and electricity Lattice structures of metals Metals are arranged in very compact and orderly patterns. Metal atoms form lattices in characteristic patterns. Malleability of Metals Metals are malleable When a force is applies the metal kernels move to a new location, changing the shape of a metal. Malleable and Ductile + + + + Force + + + + + + + +  Malleable and Ductile  Mobile e-’s allow atoms to slide by  like ball bearings in oil. + + + + Force + + + + + + + +  High Melting Point The generally high melting points indicate that metallic bonding is quite strong. Melting points increase with an increase in the number of valence electrons to the sea, since there is a greater attractive force between the cations and the electrons.  High Density Most metals have relatively high densities because metallic lattices are close-packed. Conductivity of Metals- ( Heat and electricity Good conductors of heat Metals conduct electricity because of the free moving electrons Electricity is energy due to the flow of electrons around the cations as each electron has a charge  Conduction of Heat Electrons are able to gain kinetic energy in hotter areas of the metal and are able to quickly transfer it to other parts of the metal lattice because of their freedom of movement. Heat causes the electrons to move faster and the ‘bumping’ of these electrons with each other and the protons transfers the heat. Ionic Metals are Compounds malleable are brittle. High melting Good points and conductor of Good conductor boiling points electricity of electricity (in (solid and solution liquid) and when melted) Alloys Alloys are mixtures composed of two or more elements, at least of which is a metal. Alloys are usually stronger than their individual components Steel: iron and carbon, boron, chromium, manganese, molybdenum, nickel, tungsten, vanadium. http://www.drkstreet.com/resources/metallic-bonding-animation.swf http://www.ausetute.com.au/metallic.html http://www.bbc.co.uk/schools/gcsebitesize/science/add_gateway_pre _2011/periodictable/metalsrev1.shtml http://www.educationscotland.gov.uk/highersciences/chemistry/ani mations/bondingstructure.asp The ability to conduct electricity in a solid state is a characteristic of metallic bonding. This characteristic is best explained by  A. high ionization energies  B. high electronegativity's  C. mobile electrons  D. mobile protons Which element consists of positive ions immersed in a “sea” of mobile electrons?  A. sulfur  B. nitrogen  C. calcium  D. chlorine Structure of Metals and the Electrostatic force How does the electrostatic force arise? Electrostatic forces of attraction between the positively charged cations and the negatively charged electrons hold the lattice together. A metal is therefore a seen as a rigid framework of cations immersed in a ‘sea’ of electrons that serve as the cement holding the three-dimensional cationic network together – Metallic bonding.