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Chemical Changes and Structure This unit covers (you get notes for this): • How chemist can control the rate of chemical reactions, and the enthalpy changes that take place. • Trends in the periodic table • The relationship between the arrangement of elements in the periodic table and their bonding, structure and properties. • Polar covalent bonds in the context of the bonding continuum, followed before studying intermolecular forces. Chemical changes and Structure From previous work you should know and understand the following: • Factors that affect the rate of a reaction and average rate calculations. • Atomic structure • Electron orbital's or energy levels • Valency •Covalent and ionic bonding •Physical properties of substances. Controlling the rate Overview Learn to explain how a number of key factors can influence reaction rate, using the collision theory. Factors Affecting Rate of Reaction 4 factors affect the speed of a reaction: 1. Particle size 2. Concentration of solution 3. Temperature 4. Catalysts Collision Theory a) Collision theory Learning intention Learn how chemists control reaction rates by careful consideration of the influence of concentration, temperature, surface area and collision geometry. Rates of Reaction Reactions happen at different rates. Industry needs to control reaction rates to increase production and get a good return for the investment Rates may need to be controlled for safety, or to keep the rate of production within the limit of the plant Collision Theory For a chemical reaction to occur, reactant molecules must collide. The collision must provide enough energy to break the bonds in the reactant molecules Then new chemical bonds form to make product molecules. Progress of a Reaction Reactions can be followed by measuring changes in concentration, mass and volume of reactants and products. A. Where is the reaction the quickest? B. Why does the graph level off? Rate No more products formed. C. Why does the graph curve? A The concentration of the reactants decrease with time. time 1.1 Effect of Surface Area Effect of Surface Area Particle size, the smaller the particles, the greater the surface area, the greater the chance of successful collisions. 4X4= 16 cm2 16x6=96 cm2 2x2 = 4 cm2 4X6= 24 cm2 24X8= 192 cm2 Rate and Particle Size Only the particles on the surface of a solid can be involved in a collision Higher Chemistry Eric Alan and John Harris Crushing a solid increases the surface area more particles are available for collision therefore increased rate of the reaction Effect of Surface Area Hydrochloric acid reacts with marble chips (calcium carbonate) 2HCl(aq) + CaCO3(s) CaCl2(aq) + CO2(g) + H2O (l) Rates of reaction The rate of reaction can be followed by measuring changes in Concentration Mass Volume of gas produced Measuring reaction rates Change in mass (g) Products Reactants time (s) Average rate of reaction = change in mass of product or reactant time interval Units g s-1 Measuring reaction rates Change in volume (cm3) Products Reactants time (s) Average rate change in volume of product or reactant of reaction = in time for the change to occur Units cm3 s-1 Measuring reaction rates Change in concentration (mol l-1) Products Reactants Time (s) Average rate change in concentration of product or reactant of reaction = time interval Units mol l-1 s-1 How can we follow the reaction? A gas is produced. What will happen to the gas if there is no lid on the container? What will happen to the mass? How can we follow the rate? What to do: You can follow the rate of a reaction involving gases by; 1.Measuring the volume of gases produced over time 2. Measuring the loss of mass over time How can we follow the reaction? • If we use a container fitted with a delivery tube we could measure the amount of gas produced. How? Measuring rate of reaction Two common ways: 1) Measure how fast the products are formed. • method can accurately measure volume of gas produced. 2) Measure how fast the reactants are used up by measuring the mass lost. Following the rate of a reaction–Activity 1.1 Follow the instructions on the experiment card Write down the aim, next to your note. Complete the experiment , recording your results and write a conclusion – complete all questions on the back of the instructions. What to do Record your results in a the table. Plot a graph of volume vs time using the same axes for both sets of data rate = change in volume ( the unit is cm3 s-1) time interval Swap results Each group should have a sets of results which can be used to plot graphs. 2.5 2 Loss in mass (g) 1.5 lumps powder 1 0.5 0 0 20 40 60 80 100 Time (s) rate = change in mass = ____________________ g s-1 time interval Rate over 1st 25 seconds (g s-1) rate over 2nd 25seconds (g s-1) Whole chips (C) 0.8-2 25-0 =0.05 0.35 -0.8 50-25 =0.018 Ground chips (G) 0.3-2 25-0 =0.068 0.25-0.3 50-25 =1x10-3 Now do Question 1-5 on back of instructions. Assessment Activity 1.1 Marble Chips and Acid 1. This experiment used changes in mass to follow the rate of the reaction. Suggest other methods of monitoring a chemical reaction and thus determining its rate. 2. Write a balanced equation for the reaction between marble chips (calcium carbonate) and dilute hydrochloric acid. 3. Describe the way the course of this reaction was followed. 4. Why does the curve level off after some time? 5. From your small chips graph, estimate the time it takes for - all the acid to react, - half the acid to react. 6. Explain why the time taken for all the acid to react is not exactly twice the time taken for half the acid to react. Effect of Concentration Rate and Concentration For a reaction to take place the particles must collide. Increasing the concentration of a solution increases the number of particles in the same volume. Therefore more chance of collision i.e. increased rate of the reaction Effect of Concentration • The higher the concentration, the more particles in a given space, the more chance there is of successful collisions. Using Relative (relevant) Rate 1. A reaction with high concentration took 90 seconds to complete. A similar reaction with low concentration took 160 seconds to complete. Calculate the relative rate of both reactions. 2. a) Calculate the time for reaction 1 when concentration is 0.4mol/l. b) Calculate the time for reaction 2 at 40oC Effect of Concentration –the chemical clock challenge Your challenge is to create a series of solutions that will change colour in time to music. http://www.youtube.com/watch?v=rSAa iYKF0cs Effect of Concentration –the chemical clock challenge Checklist; •Aim of the experiment •Hypothesis •Method, which variables to control and change •What to measure and how to measure it •How to record your results – table of results •What graph to draw – see page 6 of notes •Make a conclusion •Evaluate – how can you improve it/reduce the errors Effect of Concentration –the chemical clock challenge The reaction is between potassium iodide (KI) and hydrogen peroxide. The iodine clock reaction changes from colourless to blue/black Effect of Concentration –the chemical clock challenge You will carry out the reaction using a series of dilutions of the iodide solution. This will be diluted by replacing some of the volume with water. Effect of concentration –the chemical clock challenge 2I- (aq) + H2O2 (aq) I2 (aq) + 2S2O32- + 2H+ (aa) (aq) 2H2O (l) + I2 2I- (aq) + S4O62- (aq) (ag) The reaction mixture stays colourless as the iodine molecules are converted back to iodide molecules by the thiosulphate ions. Once all the thiosulphate ions have been used, a blue black colour appears suddenly as iodine reacts with starch. Relative Rate = t is the time…. 1 t Units s-1 This is a measure of how long it takes for the blue/black colour to form. (when excess I2 forms) Effect of concentration –the chemical clock challenge 1) Using syringes measure out 10cm3 sulphuric acid 0.1moll-1 10cm3 sodium thiosulphate 0.005moll-1 1cm3 starch solution 25cm3 potassium iodide solution 0.1mol l-1 Into a dry 100cm3 beaker 2) Measure out 5cm3 of hydrogen peroxide 0.1moll-1 into a syringe. Add it to the mixture as quickly as possible and start the timer. 3) Stop the clock when the mixture suddenly turns dark blue. 4) Repeat, using 20 cm3 of potassium iodide solution and 5cm3 of water with, then using repeated dilutions Effect of Concentration –the chemical clock challenge Volume Volume Concentra of water of 0.5 tion of KI (cm3) mol l-1 KI (aq) (cm3) O.0 25.0 0.1 5.0 20.00 10.0 15.0 15.0 10.0 20.0 5.0 0.06 0.02 1st Time (s) 2nd time Average Relative time Rate (1/t) Effect of Concentration –the chemical clock challenge RESULTS - Plot a graph showing the volume of potassium iodide x axis and the rate of reaction on the y axis. Effect of Changing Concerntration on Rate of Reaction Write-up 1. 2. 3. 4. 5. 6. Aim Hypothesis The independent variable (what you changed), The dependent variable (what you measured), Safety The method mentioning all the equipment used and measurements made, readings and variable kept constant/changed etc 7. A table (with headings) of your measurements, and a sample average and rate = 1/t calculation) and your line graph. 8. Your conclusion (what you found out – must mention results and link to the aim) 9. An evaluation (Comment on how well you felt exp. Went and how you can improve your investigation) Effect of Concentration - Lemonade challenge • Your task is to make up a solution containing the same concentration of sugar as a can of lemonade which contains 24g of sucrose C12 H22 O11 Effect of Concentration - Lemonade challenge Work out the mass of sucrose required to make up 100cm3 of sucrose solution of the same concentration as lemonade, assuming there are 24g in 330cm3 Make up your solution. What is the concentration of this solution? Making a standard sample • The sample is weighed and dissolved in a small volume of (deionised) water in a beaker. • The solution is transferred to a standard flask. • The beaker is rinsed and the rinsings also poured into the standard flask. • The flask is made up to the mark adding the last few drops of water using a dropping pipette. • The flask is stoppered and inverted several times to ensure thorough mixing of the solution. Effect of Concentration - Lemonade challenge Concentration (c ) is measured in moles per litre ( mol l-1) no moles = cxv 1000 To calculate the concentration you need to work out the number of mol of sugar present. Effect of Concentration - Lemonade challenge No moles = mass (g) GFM GFM = gram formula mass Effect of Concentration - Lemonade challenge • Sucrose is a non-reducing sugar – it does not react with Benedicts unless it is first hydrolysed. • Boil 10 cm3 sugar solution with 5cm3 1mol/l HCl. Then neutralise the solution with 5cm3 1mol/l NaOH. • Allow this solution to cool to room temperature. • Repeat with the lemonade solution. • Add 2cm3 of Benedicts to each sample. • Prepare a beaker of boiling water. Effect of Concentration - Lemonade challenge Add the test tubes to the boiling water If you have made up the solutions correctly, all your solutions should take the same time to change colour. Effect of Temperature on Rate of Reaction Effect of Temperature -the vanishing cross Sodium thiosulfate solution is reacted with acid. A precipitate of sulfur forms. The time taken for a certain amount of sulfur to form is used to indicate the rate of the reaction. Effect of Temperature -the vanishing cross Effect of Temperature -the vanishing cross Effect of Temperature -the vanishing cross Results Temperature (0C) Reaction time in seconds 1/time 19 32 38 51 60 105 46 36 18 12 0.0095 0.0217 0.0278 0.0556 0.0833 Effect of Temperature on Rate of Reaction 0.1 0.09 A 10 oC rise in temperature will approximately double the rate of the reaction 0.08 0.07 1/Time 0.06 0.05 0.04 0.03 0.02 0.01 0 0 10 20 30 Temperature 40 50 60 70 Temperature and Energy Effect of Temperature -the vanishing cross • The experiment can be viewed at • http://media.rsc.org/Classic%20Chem%20e xperiments/CCE-64.pdf Activation energy and reaction pathway Products Reactants 1. Before collision Reactants Energy changes 2. In full collision Activated complex Partially broken reactant bonds and partially formed product bonds 3. After collision Products If the reactants have enough combined K.E. to overcome Ea , their K.E. is converted into the energy needed to form the activation complex. Temperature and kinetic energy What do we mean by temperature and heat (thermal energy) ? The thermal energy of a system is a measure of both the potential and kinetic energy within the system. The temperature is a measure of how ‘hot’ a system is. The temperature is a measure of the average kinetic energy in a system. Temperature and Energy Energy Distribution No of molecules Ea Number of collisions which result in new products being formed. Kinetic energy Total number of collisions with sufficient K.E. energy is the area under the graph to the right of the Ea . No of molecules Temperature and Activation energy Increasing the temperature means a greater Activation EA number of molecules have energies in excessenergy of At the higher temperature EA. greyof area T2 theThe number particles T1 T2 represents the of E with energy in excess A number particles is greater. Theofblue area now witha energy represents particles with Even a small rise in temperature causes largein excess ofinthe EA. These energy excess of EA. increase in the number of particles with energy particles have enough above EA. Therefore a greater proportion energyof to cause successful collisions collisions will be successful. Energy associated with molecules The blue area is larger than the grey so at temperature T2 more particles have enough energy to cause successful collisions. Energy Distribution-Concentration No of molecules Ea Ea does not change but the number of successful collisions increases significantly, so rate increases. Increased temperature Kinetic energy No of molecules Ea Increased concentration Ea does not change but the number of successful collisions increases. Kinetic energy b) Reaction Profiles Learning intention • Learn how a potential energy diagram can be used to describe a reaction pathway, and to display activation energy and reaction enthalpy. • Learn that the activated complex is the unstable intermediate formed at the peak of the potential energy diagram. Reaction Profiles – Enthalpy (clip) Thermochemistry is the study of heat energy taken in or given out in chemical reactions. This heat, absorbed or released, can be related to the internal energy of the substances involved. Such internal energy is called ENTHALPY, symbol H. As it is only possible to measure the change in enthalpy, the symbol H, is used. H = Hp - Hr Enthalpy (products) – Enthalpy (reactants) Units kJ, kilojoules Reaction profiles -Exothermic and Endothermic Reactions Step 1: Energy must be SUPPLIED to break bonds: Step 2: Energy is RELEASED when new bonds are made: A reaction is EXOTHERMIC if more energy is RELEASED then SUPPLIED. If more energy is SUPPLIED then is RELEASED then the reaction is ENDOTHERMIC Reaction profiles -Exothermic and Endothermic Reactions -H Enthalpy of reactants Enthalpy of products +H Enthalpy of reactants Enthalpy of products Exothermic reactions give out thermal energy so the enthalpy of the products is less than that of the reactants. Endothermic reactions take in thermal energy from their surroundings. The enthaply of the products is greater than that of the reactants Reaction profiles -Exothermic and Endothermic Reactions Exothermic reactions give out heat, causing a rise in the temperature Endothermic reactions take in heat The energy change in a reaction can be shown in a potential energy diagram or reaction profile -H Path of reaction Exothermic reactions give out thermal energy H = -ve PE kJmol-1 PE kJmol-1 Reaction profiles -Exothermic and Endothermic Reactions +H Path of reaction Endothermic reactions take in thermal energy from their surroundings. H = +ve Reaction profiles -Exothermic and Endothermic Reactions A. Combustion of methane CH4 H (g) + 2O2 CH4 CO2 (g) + 2 H20 products B. Cracking of ethane C2H4 (g) + products H kJmol-1 C2H6 (g) reactants + 2O2 (g) → CO2 (g) + 2 H20 (l) (g) H negative, exothermic reaction reactants kJmol-1 (g) C2H6 (l) (g) → C2H4 (g) + H2(g) H2(g) H positive, endothermic reaction Reaction profiles -Exothermic and Endothermic Reactions Mix the following pairs of chemicals in a polystyrene cup to discover if the reactions are exothermic or endothermic Reaction 10cm3 NaOH + 10cm3 HCl 10cm3 NaHCO3 + 4 spatulas citric acid 10cm3 CuSO4 + spatula of Zn powder 10cm3 H2SO4 + Mg ribbon Temp before mixing/oC Temp after Endothermic mixing/oC or exothermic Use of the thermite reaction This reaction is used to weld railway lines together http://www.youtube.com/watch?v=vCqG 3rWtNbc&feature=related An endothermic reaction Higher Chemistry Eric Alan and John Harris Enthalpy changes and industrial processes For industrial processes it is essential that chemists can predict the quantity of heat taken in or given out. Exothermic reactions lower the temperature, slowing the reaction rate Heat must be supplied to maintain the rate of reaction – this is an expense Enthalpy changes and industrial processes Exothermic processes produce heat Heat may need to be removed to prevent the reactions proceeding beyond the capacity of the plant c) Temperature and kinetic energy Learning intention Learn to describe the relationship between temperature, kinetic energy and activation energy. Activation energy and reaction pathway Potential energy diagrams give useful information about the energy profile of a reaction. The activation energy is the minimum kinetic energy required by colliding molecules for a reaction to occur. In the diagrams shown above the activation energy appears like a ‘energy barrier’ which reactants must get over to become products. Higher Chemistry Eric Alan and John Harris Activation energy and reaction pathway 2 Breaking bonds 1 Making bonds P.E. 3 Activation Energy EA 1 2 Activated complex 3 Reaction Path Activation Energy is the additional P.E. which has to be attained by colliding molecules to form an activated complex. Activated complex is the unstable arrangement of atoms formed at the maximum of the potential energy barrier. Activation energy and reaction pathway As a reaction proceeds from reactants to products, an intermediate stage is reached at the top of the activation barrier at which a highly energetic species called an activated complex is formed. A+B → X → C+D Higher Chemistry Eric Alan and John Harris Activation energy and reaction pathway A+B → X → C+D This unstable activated complex only exist for a short period of time. From the peak of the energy barrier it can lose energy in one of two ways i.e. to the stable products or to form the reactants again. The higher the Ea the higher the barrier and the slower the reaction. Higher Chemistry Eric Alan and John Harris Ea -H Exothermic reactions give out thermal energy H = -ve Enthalpy Enthalpy Activation energy and reaction pathway Ea +H Endothermic reactions take in thermal energy from their surroundings. H = +ve Activation energy and reaction pathway 1. Mark Ea and ∆H on the PE diagrams and then calculate the value of each for the forward reaction. Ea Ea Ea ∆H ∆H ∆H A Ea = 50 KJmol-1 ∆H = -10 kJmol-1 B Ea = 30 KJmol-1 C Ea = 40 KJmol-1 ∆H = -40 kJmol-1 ∆H = +20 kJmol-1 Higher Chemistry Eric Alan and John Harris Activation energy and reaction pathway 2. Mark Ea and ∆H on the PE diagrams and then calculate the value of each for the reverse reaction. Ea Ea Ea ∆H ∆H ∆H A Ea = 60 KJmol-1 B Ea = 70 KJmol-1 C Ea = 20 KJmol-1 ∆H = +10 kJmol-1 ∆H = +40 kJmol-1 ∆H = -20 kJmol-1 Higher Chemistry Eric Alan and John Harris Activation energy and reaction pathway Ea Calculate Ea for the forward reaction Ea = 210 – 20 = 190kJ Higher Chemistry Eric Alan and John Harris d) Catalysts Learning intention Learn how a catalyst speeds up reaction rate by lowering the activation energy, and how to represent this on a potential energy diagram. Catalysts at Work Heterogeneous When the catalyst and reactants are in different States you have ‘Heterogeneous Catalysis’. They work by the adsorption of reactant molecules. E.g. Ostwald Process (Pt) for making nitric acid and the Haber Process (Fe) for making ammonia and the Contact Process (Pt) for making Sulphuric Acid. Homogeneous When the catalyst and reactants are in the same state you have ‘Homogeneous Catalysis’. E.g. making ethanoic acid from methanol and CO using a soluble iridium complex. Enzymes Are biological catalysts, and are protein molecules that work by homogeneous catalysis. E.g. invertase and lactase. Enzymes are used in many industrial processes How a heterogenous catalyst works Heterogenous Catalysis are thought to work in three stages... Adsorption Reaction Desorption Higher Chemistry Eric Alan and John Harris How a heterogenous catalyst works For an explanation of what happens click on the numbers in turn, starting with How a heterogenous catalyst works Adsorption (STEP 1) Incoming species lands on an active site and forms bonds with the catalyst. It may use some of the bonding electrons in the molecules thus weakening them and making a subsequent reaction easier. How a heterogenous catalyst works Adsorption (STEP 1) Incoming species lands on an active site and forms bonds with the catalyst. It may use some of the bonding electrons in the molecules thus weakening them and making a subsequent reaction easier. Reaction (STEPS 2 and 3) Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur. This increases the chances of favourable collisions taking place. How a heterogenous catalyst works Adsorption (STEP 1) Incoming species lands on an active site and forms bonds with the catalyst. It may use some of the bonding electrons in the molecules thus weakening them and making a subsequent reaction easier. Reaction (STEPS 2 and 3) Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur. This increases the chances of favourable collisions taking place. Desorption (STEP 4) There is a re-arrangement of electrons and the products are then released from the active sites Examples of heterogenous catalysts Format Metals Ni, Pt Fe Rh, Pd hydrogenation reactions Haber Process catalytic converters Oxides Al2O3 V2O5 dehydration reactions Contact Process FINELY DIVIDED increases the surface area provides more collision sites IN A SUPPORT MEDIUM maximises surface area and reduces costs How a heterogenous catalyst works In some cases the choice of catalyst can influence the products Ethanol undergoes different reactions depending on the metal used as the catalyst. The distance between active sites and their similarity with the length of bonds determines the method of adsorption and affects which bonds are weakened. Copper Dehydrogenation (oxidation) C2H5OH ——> CH3CHO + H2 Alumina Dehydration C2H5OH ——> C2H4 + H2O How a heterogenous catalyst works Poisoning Impurities in a reaction mixture can also adsorb onto the surface of a catalyst thus removing potential sites for gas molecules and decreasing efficiency. expensive because... the catalyst has to be replaced the process has to be shut down examples Sulphur poisons iron in Haber process Lead poisons Pt in catalytic converters in cars Investigating Catalysis Decomposition of hydrogen peroxide Hydrogen → water + oxygen 2H2O2 → 2H2O + O2 What will you see during the reaction? What is the test for oxygen? Which type of catalysis? Did the experiment involve heterogeneous or homogeneous catalysis? An Example of a homogenous catalyst Higher Chemistry Eric Alan and John Harris Catalysts and Potential energy diagrams Catalysts work by providing… “AN ALTERNATIVE REACTION PATHWAY WHICH HAS A LOWER ACTIVATION ENERGY” Potential energy graphs and catalysts Uncatalysed 75 - 60 50 - Catalysed Reactants P.E. Products 25 - Reaction path Potential energy graphs and catalysts Activation energy Ea for the forward uncatalysed reaction 75 - 60 - Activation energy Eafor the forward catalysed reaction Reactants 50 - P.E. Products 25 - Reaction path Catalysts lower the activation energy needed for a successful collision. Potential energy graphs and catalysts Activation energy Ea for the reverse uncatalysed reaction reaction 75 - Activation energy Ea for the reverse catalysed reaction 60 50 - Reactants P.E. Products 25 - Reaction path Catalysts lower the activation energy needed for a successful collision. Potential energy graphs and catalysts 75 - Catalysts lower the activation energy needed for a successful collision. 60 50 - Reactants P.E. Products 25 - Reaction path ∆H Effect of catalyst – forward reaction Effect of catalyst – reverse reaction Activation energy No change Lowered No change Lowered Catalysts A catalyst speeds up the reaction by lowering the activation energy. A catalyst does not effect the enthalpy change for a reaction A catalyst speeds up the reaction in both directions and therefore does not alter the position of equilibrium or the yield of product, but does decrease the time taken to reach equilibrium. Energy distribution and catalysts Ea No of Collisions with a given K.E. Un-catalysed reaction Kinetic energy Ea Total number of collisions (area under the graph) with sufficient K.E. energy to create new products. Catalysed reaction Ea is reduced Concentrated solutions of hydrogen peroxide are used in the propulsion systems of torpedoes. Hydrogen peroxide decomposes naturally to form water and oxygen: 2H2O2(aq) → 2H2O(ℓ) + O2(g) ΔH = −196∙4 kJ mol–1 Transition metal oxides act as catalysts in the decomposition of the hydrogen peroxide. Unfortunately, there are hazards associated with the use of hydrogen peroxide as a fuel in torpedoes. It is possible that a leak of hydrogen peroxide solution from a rusty torpedo may trigger an explosion. Using your knowledge of chemistry, comment on why this could happen. Trends in the periodic table and bonding Overview This section studies • how the elements are arranged in the Periodic Table, • the structure and bonding of the first twenty elements, • Helps you explain key periodic trends in physical properties and relate these to the bonding continuum. a) The arrangement of elements in the periodic table Learning intention • Learn how the elements are organised into groups and periods in order of increasing atomic number • To identify important classifications of elements within the periodic table. It has taken many years of work by many scientists to find out about the elements that we know about now (and there may be more that we don’t know about yet). Ancient Times - Prior to 1 A.D. Gold Silver Copper Iron Mercury Carbon Sulphur LeadTin Timeline of the elements Arsenic (Magnus ~1250) Antimony (17th century or earlier) Phosphorus (Brand 1669) Zinc (13th Century India) Cobalt (Brandt ~1735) Platinum (Ulloa 1735) Nickel (Cronstedt 1751) Bismuth (Geoffroy 1753) Robert Boyle In 1661, Robert Boyle showed that there were more than just four elements as the ancients had assumed. Boyle defined an element as a pure substance that cannot be decomposed into any simpler substance. Time line of elements Hydrogen (Cavendish 1766) Nitrogen (Rutherford 1772) Oxygen (Priestley; Scheele 1774) Chlorine (Scheele 1774) Manganese (Gahn, Scheele, & Bergman 1774) Molybdenum (Scheele 1778) Tungsten (J. and F. d'Elhuyar 1783) Tellurium (von Reichenstein 1782) Lavoisier 1789 http://web.bilkent.edu.tr/ The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light. Introduced a logical system for naming compounds and helped introduce the metric system Time of the chemists Uranium (Peligot 1841) Strontium (Davey 1808) Titanium (Gregor 1791) Yttrium (Gadolin 1794) Vanadium (del Rio 1801) Chromium (Vauquelin 1797) Beryllium (Vauquelin 1798) Niobium (Hatchett 1801) Tantalum (Ekeberg 1802) Atomic Weights John Dalton, 1803, was the first chemist to use the term ‘atom’ He used this idea to explain how elements react together to form molecules. Dalton suggested that it should be possible to compare the masses of atoms. Hydrogen 1 Carbone 4.2 Oxygen 5.5 Water 6.5 Sulphur 14.4 Sulphuric Acid 25.4 www.bioanalytical.com Cerium (Berzelius & Hisinger; Klaproth 1803) Palladium (Wollaston 1803) Rhodium (Wollaston 18031804) Osmium (Tennant 1803) Iridium (Tennant 1803) Sodium (Davy 1807) Potassium (Davy 1807) Barium (Davy 1808) Calcium (Davy 1808) Magnesium (Black 1775; Davy 1808) Boron (Davy; Gay-Lussac & Thenard 1808) Iodine (Courtois 1811) Lithium (Arfvedson 1817) Cadmium (Stromeyer 1817) Selenium (Berzelius 1817) Silicon (Berzelius 1824) Zirconium (Klaproth 1789; Berzelius 1824) Aluminum (Wohler 1827) Bromine (Balard 1826) Thorium (Berzelius 1828) Lanthanum (Mosander 1839) Terbium (Mosander 1843) Erbium (Mosander 1842 or 1843) Ruthenium (Klaus 1844) Cesium (Bunsen & Kirchoff 1860) Rubidium (Bunsen & Kirchoff 1861) Thallium (Crookes 1861) Indium (Riech & Richter 1863 Newlands Newlands in 1865, placed elements in order of www.chemsoc.org succession of atomic weights noticed a pattern, noticed that the 8th one was a ‘kind of repetition of the 1st. He called this the ‘Law of Octaves’. Element Atomic weights Element Atomic Weights Element Atomic Weights Hydrogen 1 Fluorine 8 Chlorine 15 Lithium 2 Sodium 9 Potassium 16 Beryllium 3 Magnesium 10 Calcium 17 Boron 4 Aluminium 11 Chromium 18 Carbon 5 Silicon 12 Titanium 19 Nitrogen 6 Phosphorus 13 Manganese 20 Oxygen 7 Sulphur 14 Iron 21 Lothar Meyer www.apsidium.com Meyer in 1869, independently, put forward a similar list of elements. Meyer plotted graphs of melting point, boiling point and atomic volume against atomic mass. He found the properties varied in a regular way i.e. periodically Mendeleev (1869) In 1869 he published ‘Principles of Chemistry’ - proposed the modern Periodic Table. elmoscow.ru - elements with similar properties were placed together - he left gaps for new 'undiscovered' elements. - he predicted properties of undiscovered elements - arranged in order of increasing relative atomic mass - some elements were out of order therefore modern table is arranged in Atomic Number Meyer recognised Mendeleev’s work and both where awarded The Davy medal for Chemistry in 1882. The world’s first view of Mendeleev’s Periodic Table – an extract from Zeitschrift fϋr Chemie, 1869. Correct predictions The greatness of Mendeleev was that not only did he leave spaces for elements that were not yet discovered but he predicted properties of five of these elements and their compounds including gallium which he called eka-aluminium. In Paris (1875) Paul Emile Lecoq de Boisbaudran discovered an element he named gallium after the Latin name for France Eka-aluminium (Ea) Gallium (Ga) Atomic weight About 68 69.72 Density of solid 6.0 g/cm3 5.9 g/cm3 Melting point low 29.78oC Valency 3 3 Method of discovery Probably from its spectrum Spectroscopically Oxide Formula Ea2O3, density 5.5 g/cm3. Soluble in both acids and alkalis. Formula Ga2O3, density 5.88 g/cm3. Soluble in both acids and alkalis. Sir William Ramsay One thing that Mendeleev did not predict was the discovery of a whole new Group of elements, the noble gases, discovered by Scot William Ramsay and th co-workers during the last decade of the 19 century. Groups The Periodic Table Periods Groups - vertical columns. - elements in a group show specific similarities. - common names : Alkali Metals, Halogens, Noble Gases. - increasingly metallic down a group, nonmetallic up a group. - outer shell electrons determine the group number. Periods - horizontal rows - two short periods, four long periods. - elements change from metallic to nonmetallic across a period. - number of the shell determines the period. Periodicity It is the ideas of Meyer and Mendeleev that we will make use of to try and understand the relationships between the first 20 elements. Periodicity - As you move across a Period from left to right, a pattern emerges. A similar pattern appears on crossing the next Period. •Properties which behave in this way are said to be periodic. Periodicity is the regular recurrence of similar element properties. • Density - Lothar Meyer Curve Variation of density (g cm-3) with atomic number Adapted from New Higher Chemistry E Allan J Harris period 2 (Li - Ne) maximum at boron (B) group3 period 3 (Na - Ar) maximum at Aluminium (Al)- group 3 Variation of density (g cm-3) with atomic number Adapted from New Higher Chemistry E Allan J Harris In general in any period of the table, density first increases from group 1 to a maximum in the centre of the period, and then decreases again towards group 0 5th 4th 2nd 3rd Variation of density (g cm-3) with atomic number Adapted from New Higher Chemistry E Allan J Harris down a group gives an overall increase in density Melting point - Lothar Meyer Curve Variation of melting point with atomic number Adapted from New Higher Chemistry E Allan J Harris Determined by the strength of intermolecular bonding, between particles period 2, peak at carbon period 3, peak at silicon In general the forces of attraction (intermolecular bonding) for elements on the left of the table must be stronger, or more extensive than between the particles on the right. Variation of melting point with atomic number Adapted from New Higher Chemistry E Allan J Harris Down group 1 the alkali metals m.pt. decrease there must be a decrease in the force of attraction between the particles Variation of melting point with atomic number Adapted from New Higher Chemistry E Allan J Harris Down group 7 the halogens m.pt. increases there must be a increase in the force of attraction between the particles Boiling point - Lothar Meyer Curve Variation of boiling point with atomic number Adapted from New Higher Chemistry E Allan J Harris period 2, peak at carbon period 3, peak at silicon In general we see the same trend in boiling point across the period Variation of boiling point with atomic number Adapted from New Higher Chemistry E Allan J Harris Down group 1 the alkali metals b.p. decrease once again there must be a decrease in the force of attraction between the particles Variation of boiling point with atomic number Adapted from New Higher Chemistry E Allan J Harris Down group 7 the halogens b.p. increases once again there must be a increase in the force of attraction between the particles b) Periodic trends in covalent radii and ionisation energies Learning intention • Learn the definitions of covalent radius and first ionisation energy • find out how to explain their periodic trends in relation to atomic size, nuclear charge and the screening effect of inner shells of electrons. Covalent radius Atomic Size There is no definite edge to an atom. However, bond lengths can be worked out. Covalent radius, ½ the distance between nuclei. To find the bond length, add 2 covalent radii together. pm = picometre X 10 – 12 m Variation of covalent radius with atomic number Adapted from New Higher Chemistry E Allan J Harris The covalent radii of the elements in any period decrease with increasing atomic number. Variation of covalent radius with atomic number Adapted from New Higher Chemistry E Allan J Harris The covalent radii of the elements in any group increase with increasing atomic number. Variation of covalent radius with atomic number Adapted from New Higher Chemistry E Allan J Harris No values are given for the Nobel gases Why? Unreactive so do not form bonds Covalent radius COVALENT RADIUS Explain the change in covalent radius as you go along a period. The covalent radius of an element is half the distance between the nuclei of two of its covalently bonded atoms. The covalent radius decreases as you go along a period As you go along a period there is a greater positive charge on the nucleus The shells or energy levels of electrons are more strongly attracted to the nucleus and therefore the size of the atoms decreases. COVALENT RADIUS Explain the change in covalent radius as you go down a group. The covalent radius of an element is half the distance between the nuclei of two of its covalently bonded atoms. The covalent radius increases as you go down a group. As you go down a group there are more energy levels of electrons. The outer electrons are further away from the nucleus so the atoms are larger. Trend in Ionisation energy Ionisation energies This is defined as "the amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state” Energy e e M (g) M+(g) + e 1st ionisation + M +(g) The outermost electron will be the most weakly held and is removed first Ionisation energies This is defined as "the amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state” Energy e M (g) M+(g) + e 1st ionisation e + M 2+ (g) M(g)+ M(g)2+ + e 2nd ionisation The ionisation energy is an enthalpy change and therefore is measured per mole. Units kJmol-1 (kilojoules per mole). Ionisation energies kJ mol-1 Overall increase along period Decrease down group Li 526 Be 905 B 807 C N O F Ne 1090 1410 1320 1690 2090 Na 502 K 425 Rb 409 Mg 744 Ca 596 Sr 556 Al 584 Ga 577 In 556 Si 792 Ge 762 Sn 715 P 1020 As 953 Sb 816 S 1010 Se 941 Te 870 Cl 1260 Br 1150 I 1020 Ar 1530 Kr 1350 Xe 1170 FIRST IONISATION ENERGY Explain the change in first ionisation energy as you go down a group. The first ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state. The first ionisation energy decreases as you go down a group. As you go down a group there are more energy levels of electrons. The outer electron is further from the nucleus. The inner electrons shield the outer electron from the effect of the nucleus. Less energy is needed to remove the outer electron. First and Second ionisation energies of the first 20 elements Adapted from New Higher Chemistry E Allan J Harris Down a group first ionisation energy decreases FIRST IONISATION ENERGY Explain the change in first ionisation energy as you go along a period. The first ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state. The first ionisation energy increases as you go along a period As you go along a period there is a greater positive charge on the nucleus. There is a greater attraction between the outer electron and the nucleus. More energy needs to be supplied to remove the electron. First and Second ionisation energies of the first 20 elements Adapted from New Higher Chemistry E Allan J Harris In each period there is an overall increase peaking at the noble gas First and Second ionisation energies of the first 20 elements Adapted from New Higher Chemistry E Allan J Harris For each element the second ionisation energy is higher than the first ionisation energy. First and Second ionisation energies of the first 20 elements Adapted from New Higher Chemistry E Allan J Harris It is worth noting the Nobel gases have the highest value for each period. This goes some way to explaining the great stability of filled orbital's and the resistance of the Nobel gases to form compounds. Successive ionisation Energies first ionisation energy E(g) second ionisation energy E+(g) third ionisation energy E 2+(g) fourth ionisation energy E 3+(g) E+(g) + eE 2+ (g) + eE 3+ (g) + eE 4+ (g) + e- - ionisation energies increase as successive electrons are removed - removing an electron from a filled inner shell requires a large increase in energy The first four ionisation energies of aluminium, for example, are given by Al(g) Al+(g) Al2+(g) Al+(g) + e- 1st I.E. = 577 kJ mol-1 Al2+ (g) + e- 2nd I.E. = 1820 kJ mol-1 Al3+ (g) 3rd I.E. = 2740 kJ mol-1 + e- In order to form an Al3+(g) ion from Al(g) you would have to supply: 577 + 1820 + 2740 = 5137 kJ mol-1 Explain why the second ionisation energy of an element is always greater than the first ionisation energy: First ionisation energy – first mole of electrons removed Second ionisation energy – second mole of electrons removed M(g) M+(g) + e M+(g) M2+(g) + e In the second ionisation energy negative electrons are being removed from positive ions rather than neutral atoms. In the positive ion there is a greater attraction for the electrons so more energy is needed to remove the second mole of electrons. Explain why the second ionisation energy of K is much greater than the second ionisation energy of Mg: K (g) 2,8,8,1 K+ (g) + e 2,8,8 K+ (g) K2+ (g) + e 2,8,8 2,8,7 Mg (g) Mg+ (g) + e 2,8,2 2,8,1 Mg+ (g) Mg2+ (g) + e 2,8,1 2,8 The second ionisation of K involves removing an electron from a stable electron arrangement. The second ionisation of Mg involves removing an electron to form a stable electron arrangement. This requires a lot of energy This requires less energy Nobel gas compounds If some other change can compensate for the energy required then ionic compounds of Nobel gases can be made. Can you suggest why the first Nobel gas compound prepared contained Xe? Questions for you to try: 1. Explain why the third ionisation energy of Magnesium is so much greater than its second. 2. Calculate the energy change for the following changes. a)Ca (g) → Ca2+ (g) + 2eb)B2+(g) → B4+ (g) + 2e3. Which of the following equations represents the first ionisation energy of fluorine? A. F–(g) → F(g) + e– B. F–(g) → F2(g) + e– C. F(g) → F+(g) + e– D. F2(g) → F+(g) + e– 4. Which line in the table is likely to be correct for the element francium? State at 30 °C First ionisation energy/kJ mol–1 A solid less than 382 B liquid less than 382 C solid greater than 382 D liquid greater than 382 5. As the atomic number of the alkali metals increases A the first ionisation energy decreases B the atomic size decreases C the density decreases D the melting point increases. 6. Explain why a) A potassium atom is larger than a sodium atom b) The Chlorine atom is smaller than a sodium atom 7. Atoms of different elements are different sizes. What is the trend in atomic size across the period from sodium to argon? 8. Which of the following reactions refers to the third ionisation energy of aluminium? A Al(s) → Al3+(g) + 3e– B Al(g) → Al3+(g) + 3e– C Al2+(g) → Al3+(g) + e– D Al3+(g) → Al4+(g) + e– 9. Atoms of different elements have different ionisation energies. Explain clearly why the first ionisation energy of potassium is less than the first ionisation energy of sodium. Answers 1.Explain why the third ionisation energy of Magnesium is so much greater than its second. Removing the third mole of electrons involves breaking into a stable complete energy level of electrons. This energy level is closer to the nucleus so more attraction between the outer electron and nucleus. 2. Calculate the energy change for the following changes. a)Ca (g) → Ca2+ (g) + 2e- 596 + 1160 = 1756 kJmol-1 b)B2+(g) → B4+ (g) + 2e3660 + 25000 = 28660 kJmol-1 3. Which of the following equations represents the first ionisation energy of fluorine? A F–(g) → F(g) + e– B F–(g) → F2(g) + e– C F(g) → F+(g) + e– D F2(g) → F+(g) + e– C 4. Which line in the table is likely to be correct for the element francium? State at 30 °C First ionisation energy/kJ mol–1 A solid less than 382 B liquid less than 382 C solid greater than 382 D liquid greater than 382 B 5. As the atomic number of the alkali metals increases A the first ionisation energy decreases B the atomic size decreases C the density decreases D the melting point increases. A 6. Explain why a) A potassium atom is larger than a sodium atom Potassium has an extra energy level (shell) of electrons. b) The Chlorine atom is smaller than a sodium atom Both atoms have the same number of energy levels, but the chlorine has a greater nuclear charge than sodium. This attracts the outer electrons more strongly and causes the decrease in atomic radius. 7. Atoms of different elements are different sizes. What is the trend in atomic size across the period from sodium to argon? Decreases or gets smaller 8. Which of the following reactions refers to the third ionisation energy of aluminium? A Al(s) → Al3+(g) + 3e– B Al(g) → Al3+(g) + 3e– C Al2+(g) → Al3+(g) + e– D Al3+(g) → Al4+(g) + e– C 9. Atoms of different elements have different ionisation energies. Explain clearly why the first ionisation energy of potassium is less than the first ionisation energy of sodium. Potassium has more electron shells (or outer electron is further from the nucleus) The inner electrons (electron shells) shield (screen) the outer electron from the attraction of the nucleus. Therefore the outer electron is held less tightly in potassium Periodic Table of Data Visual database of the physical and thermochemical properties of the chemical elements which allows the user to plot graphs and tables, play games and view diagrams. Periodic Landscapes The Periodic Landscape images are computer-generated landscape views and models based on patterns and relationships within the periodic table. The models are sculpted to achieve a sense of general trends or patterns. Bonding and structure in the first 20 elements Overview Learn how the first twenty elements can be classified into groups according to their bonding and structure. Chemical Bond Intra-molecular (Metallic, ionic and covalent bonding) Inter-molecular (between molecules) Van der Waals’ http://www.educationscotland .gov.uk/highersciences/chemi stry/animations/intermolecula rforces.asp This animation describes and explains the key intermolecular forces of attraction (Van der Waals forces of attraction) including London dispersion forces, permanent dipole-permanent dipole attractions and Hydrogen Bonding. Types of bonding in elements • Metallic • Covalent • Both of these are intra-molecular forces – i.e. bonds between atoms! The Metals Metallic bonding Groups I, II, III Metallic elements Delocalised electron + + + + + + + + Positive nucleus (core) Electron shells The outer shell in metals is not full and so the outer electrons in metal atoms can move randomly between these partially filled outer shells. The electrons are delocalised (sometimes called a ‘sea’ or ‘cloud’ of Electrons) i.e. they are held in common by all the atoms. Metallic bonding is the strong electrostatic force of attraction, between the positive charged ions, formed by the loss of the outer shell electrons of a metal atom and these delocalised electrons. The positive metal ions are held together by this electron “Glue”. The outer electrons are delocalised and free to move throughout the lattice. The greater the number of electrons in the outer shell the stronger the metallic bond. So the melting point of Al>Mg>Na Conduction of Metals Metallic Character The strength of the metallic bond depends on 1) The elements tendency to lose electrons (ionise) 2) The packing arrangement of the metal atoms. 3) The size of the atom. 4) The number of valence electrons in the outer most shell. 5) The number of shells. Physical properties of metals A. Metals are malleable and ductile Metal atoms can ‘slip’ past each other because the metallic bond is not fixed and it acts in all directions. Physical properties of metals B. Conduction of electricity and thermal energy. Solid and liquid metals conduct heat and electricity. The delocalised electrons are free to move in the solid lattice. These mobile electrons can act as charge carriers in the conduction of electricity or as energy conductors in the conduction of heat. Physical properties of metals C. Change of state In general, metals have high melting and boiling points because of the strength of the metallic bond. When a metal is molten the metallic bonds are still present. B.p.’s are much higher as you need to break the metallic bonds throughout the metal lattice. Metal boiling point trends Down a group Boiling point of alkali metals 1600 Boiling point /oC 1400 1200 1000 800 Series1 600 400 200 0 Lithium Sodium Potassium Metal Boiling point across a period Boiling point /oC 3000 2500 2000 1500 Series1 1000 500 0 Potassium Calcium Metal Gallium The strength of metallic bonding decreases the forces of attraction get weaker. Across a period The covalent radius decreases as the positive core is increasing in charge, this has the effect of pulling the outer closer to the nucleus the forces of attraction increase. The Non-metals Nobel gases Group 0 Noble gases He Noble gases have full outer electron shells ++ They do not need to combine with other atoms. They are said to be monatomic. Group 0 are all gases and exist as individual atoms. However, the monatomic gases do form weak inter-atomic bonds at very low temperatures. London dispersion forces Learning intention Learn about these weak intermolecular forces of attraction – they exist between molecules of all non-metal discrete molecular elements and the atoms of monatomic noble gases. London Dispersion Forces Monatomic elements ++ Sometimes the electrons can end up on one side of the atom, i.e. the electron cloud can wobble ++ This means that one side of the atom is more negative than the other side. δ+ δ- δ- δ+ These charges are given the symbol δ ‘delta’ A temporary dipole is therefore formed. δ+ A dipole can induce other atoms to form dipoles, resulting in dipole –dipole attraction. δLondons forces Monatomic elements London dispersion forces are very weak attractive forces Noble gases b.p.’s 180 160 166 140 120 121 100 b.p / K 80 87 60 40 20 4 Helium Neon Argon Krypton Xeon 27 0 K = -293o C 0 B.p.’s increase as the size of the atom increases. This happens because the London forces increases with increasing number of electrons. The more electrons the bigger the temporary dipole, the stronger the London forces – so higher b.p. The Non-metals 1. Discrete Covalent molecules Relating physical properties to intermolecular forces in Discrete Covalent molecules Learning intention Learn how to explain differences in physical properties such as viscosity, melting point and boiling point in terms of differences in strength of intermolecular forces. Covalent molecular elements Most non-metals exist as discrete covalent molecules where their atoms are held together by strong covalent bonds. e.g. O2,S,CL2,P,N2,F2...... Discrete molecules have a definite formula with a definite number of atoms bonded together. They are very small(discrete) molecules. The discrete molecules are held by weak intermolecular forces called London Dispersion Forces. This means they have .......?..... Melting and boiling points. Covalent Bond - Revision Chlorine atom 2,8,7 17+ Chlorine molecule Cl2 2,8,8 17+ 17+ diatomic A covalent bond is formed when a pair of electrons are shared. The atoms in a covalent bond is the mutual attraction of two positive nuclei for a shared pair of electrons. This makes it very strong. Group IV Buckminster fullerene (Bucky Balls) were discovered in the 1980’s. C60 C70 Due to the large molecules , fullerenes have stronger London forces between their molecules, compared to elements made from smaller molecules. Fullerenes are a family of carbon molecules made up of rings with definite formula. They are discrete covalent molecules Group V Strong triple covalent bond N Weak Londons force N N N Strong intra-molecular bonding and weak inter-molecular bonding exist in this diatomic molecule. N 2 m.p. -210o C Phosphorus P4 m.p. 44oC Strong covalent bonds Weak Londons forces Group VI Strong double covalent bond O Weak Londons force O O O Strong intra-molecular bonding and weak inter-molecular bonding exist in this diatomic molecule. O 2 m.p. -218o C Sulphur S8 Weak Londons forces m.p. 113oC Higher m.p. because there arestronger Londons’ forcesbetween larger molecules(more electrons). Group VII Fluorine F2 Strong covalent bond F Weak London forces Strong intra-molecular bonding (Covalent) and weak inter-molecular bonding (London forces) exist in this diatomic molecule. F 2 m.p. -220o C F F F Chlorine Cl2 Cl Strong covalent bond Weak London forces Cl Cl Cl Strong intra-molecular bonding (?) and weak inter-molecular bonding (?) exist in this diatomic molecule. Cl 2 m.p. -101oC Halogens b.p.’s 500 450 457 400 b.p./ K Fluorine 350 Chlorine 300 332 250 200 238 Bromine Iodine 150 100 50 85 0 As the size of the halogen atom increases (more electrons), so does the size of the London forces between the halogen molecule. Covalent Network Elements In the first 20 elements, only Boron, Carbon and Silicon have covalent network structures. Carbon Diamond Diamond forms an infinite 3D network structure. Each carbon atom forms 4 covalent bonds to 4 other carbon atoms. Very rigid strong structure. Diamond is one of the hardest materials known to man. C sublimes 3642oC Graphite Carbon bonded to only 3 other Carbons So the spare electrons are delocalised and so free to move. Graphite is a conductor. London forces between the layers allows layers to slide over each other. Graphite can be used as a lubricant Silicon Silicon has the same infinite 3D network structure as diamond Si mp 1410oC. Trending in Properties – Revisited The link to bonding Density change across a period 3 2.5 Density g/cm3 Sodium Magnesium Aluminium Silicon Phosphorus Sulphur Chlorine Argon 2 1.5 1 0.5 0 Na Mg Al Si P S Cl Ar Na to Al the atom size decreases leading to greater packing in metal latt Si is a covalent network, tightly packed atoms in covalent lattice. P and S are covalent molecular solids with quite densely packed molecules. Cl and is a covalent molecular gas at room temperature. Ar and is a monomolecular gas at room temperature. Bonding in the first 20 elements Covalent molecular gases Li These elements occur as diatomic Covalent networks (two atom) molecules with strong Giant network of atoms withbetween strong the atoms covalent bonds covalent bonds between the atoms. (intramolecular bonds) and weak Londons Very high melting andforces boilingbetween points. the molecules (intermolecular bonds). The weak Londons forces mean low melting and boiling points. Be CC N F B O Na Mg K Ca H Metallic bonding. Giant network of positively charged nuclei surrounded by delocalised electrons. Delocalised electrons make these elements good conductors. Al Si P Cl S He Ne Ar Covalent molecular solids Polyatomic (many atom) molecules. Fullerenes C60 C70 Monatomic elements P4 and S8. These molecules have many The noblelarger gasesLondon exist as individual electrons and this produces (monatomic) atoms. There are only weak forces than the diatomic molecules. Londons forces between the atoms. The stronger London forces (temporary Very energy is needed to break dipoles) gives these twolittle elements higher these forces and so the noble gases melting and boiling points. have very low melting and boiling points. These two elements are solids at room temperature. Diagram shows part of the covalent network of carbon Uneven distribution of the electrons in atoms in diamond. the electron cloud create temporary dipoles (d atom + and is d-)covalently which result in a weak Each carbon attraction between atoms which come bonded to 4 other carbon close to each other. These weak atoms. attractions are called Londons forces. d+ dd+ d- H He Li Be B CC N O F Ne Na Mg Al Si P S S Cl Ar K Ca Metallic bonding with a network of positively charged nuclei surrounded by a ‘sea’ of delocalised electrons. Diagram Strong showscovalent S8 bonds molecules between in sulphur the atoms inside -thelondons - forces -molecules. with the diatomic shown by the dotted + + + + + + + lines. + + + + + + + These large Weakmolecules van der Waal’s have +stronger + + Londons + + molecules + + forces between forces than the which come close to each diatomic molecules. -other. - - - - http://www.educationscotland.gov.uk/highersciences/c hemistry/animations/bondingstructure.asp This interactive animation provides a visual representation of the bonding and structure of the first twenty elements in the periodic table, taking into account both the intra- and inter-molecular forces involved. Questions on elements – bonding and structure 1. Explain why the covalent network elements have high melting and boiling points. 2. Explain why the discrete molecular and monatomic elements have low melting and boiling points. 3. Does diamond conduct electricity? Explain. 4. Does graphite conduct electricity? Explain. 5. How does the hardness of diamond compare with graphite? Explain. 6. Give a use for both diamond and graphite. 7. Complete the following table: Questions on elements – bonding and structure 7. Complete the following table: Type of bonding and structure Metallic solids Properties ……………. of electricity Covalent network solids ……….. …. melting points ……………. of electricity exception ………………. Covalent molecular solids ………….. melting points …………… of electricity Covalent molecular (diatomic) gases and monatomic gases …………… boiling points b) Enthalpies of combustion Learning intention Learn the definition of enthalpy of combustion, which can be directly measured using a calorimeter. National 5 = Energy of Combustion Higher = Enthalpy(ΔH) of combustion • Definition: The enthalpy of combustion is the heat energy given out when 1 mole of fuel burns completely in oxygen. • The enthalpy of combustion of methane can be represented by the equation • CH4(g) + O2 (g) CO2(g) + H2 O(l) Enthalpy of combustion The heat energy released when alcohols burn can be measured The enthalpy of combustion of a substance is the amount of energy given out when one mole of a substance burns in excess oxygen. Calculating Enthalpy of CombustionSpecific heat capacity Calculating the energy change during a chemical reaction in water. Calculation (a)The heat energy gained by the water (Eh) can be calculated using the formula: H = c. m. T c m T =specific heat capacity =mass in Kg =temperature change The mass of water can be calculated by using the fact that 1 ml = 1 g. The value for c is usually taken as 4.18 kJ kg –1 oC-1 Measuring the enthalpy of combustion of alcohols A Procedure? Weigh a filled alcohol burner Measure 100 cm3 water into a copper calorimeter Take temperature of the water Light the burner and use it to heat the water for 3 minutes. Stir the water and take the highest temperature reached Reweigh the burner and remaining fuel Calculate the energy for mass burned Calculate the Enthalpy for 1 mole of the alcohol Enthalpy of combustion Procedure 1. Weigh the spirit burner (already containing ethanol) with its cap on and record its mass. (The cap should be kept on to cut down the loss of ethanol through evaporation) 2. Using the measuring cylinder, measure out 100 cm3 of water into the copper can. 3. Set up the apparatus as directed by your teacher/lecturer. 4. Measure and record the temperature of the water. 5.Remove the cap from the spirit burner and immediately light the burner. 6.Slowly and continuously stir the water with the thermometer. After 3 minutes, recap the spirit burner and measure and record the maximum temperature of the water. 7. Reweigh the spirit burner and record its mass. CALCULATION Suppose 0.25 g of ethanol had been burned and the temperature of 100cm3 water had risen by 12.5 °C. Part 1 The heat energy gained by the water (Eh) is calculated using the formula: Eh = c m T Eh = 4.18 x 0.10 x 12.5 = 5.225 kJ We assume that the heat energy released by the burning ethanol is gained only by the water. The heat energy released on burning 0.25 g of ethanol = 5.225 kJ However Enthalpy is for 1 mole of a substance so………. Part 2 Work out GFM (1 mole in grams) of alcohol……. Ethanol: CH3CH2OH Mass of 1 mole = 2(12) + 6(1) + 16 = 46 g We can now calculate the heat energy released on burning 1 mole of ethanol. (The Enthalpy of Ethanol) 0.25g 5.225kJ x 5.225 46g 46 =0.25 961 kJ The enthalpy of combustion of ethanol = - 961 kJ mol-1 (A negative sign is used because combustion is an exothermic reaction) Measuring the enthalpy of combustion of alcohols The heat energy gained by the water (Eh) is calculated using the formula: Eh = c m ∆T Eh = x x = kJ Measuring the enthalpy of combustion of alcohols We assume that the heat energy released by the burning alcohol is gained only by the water. The heat energy released on burning ……….. g of ……………anol So one mole .............. g of ................anol ………….. kJ ....................kJ The enthalpy of combustion of …………anol = -………….. kJ mol-1 (A negative sign is used because combustion is an exothermic reaction) Sources of inaccuracy • Heat loss to surroundings • Incomplete combustion • Possible loss of fuel by evaporation from wick Calorimetry To eliminate these inaccuracies a bomb calorimeter is used The burning fuel (or food) is supplied with oxygen to encourage complete combustion The combustion chamber is entirely surrounded so there is no heat loss to the surroundings Commercial ‘bomb’ calorimeters Databooklet Values The calorimeter is heated electrically. Energy required to heat the entire apparatus by 1 0C is calculated. Enthalpy of combustion Worked example 1. 0.19 g of methanol, CH3OH, is burned and the heat energy given out increased the temperature of 100g of water from 22oC to 32oC. Calculate the enthalpy of combustion of methanol. –704 kJ mol-1 Enthalpy of combustion Worked example 1. 0.19 g of methanol, CH3OH, is burned and the heat energy given out increased the temperature of 100g of water from 22oC to 32oC. Calculate the enthalpy of combustion of methanol. –704 kJ mol-1 Worked example 2. 0.22g of propane was used to heat 200cm3 of water at 20oC. Use the enthalpy of combustion of propane in the data book to calculate the final temperature of the water. 33.3oC Calculations for you to try. 1. 0.25g of ethanol, C2H5OH, was burned and the heat given out raised the temperature of 500 cm3 of water from 20.1oC to 23.4oC. Calculate the enthalpy of combustion of ethanol 2. 0.01 moles of methane was burned and the energy given out raised the temperature of 200cm3 of water from 18oC to 28.6oC. Calculate the enthalpy of combustion of methane. 3. 0.1g of methanol, CH3OH, was burned and the heat given out used to raise the temperature of 100 cm3 of water at 21oC. Use the enthalpy of combustion of methanol in the data booklet to calculate the final temperature of the water. 4. 0.2g of methane, CH4, was burned and the heat given out used to raise the temperature of 250 cm3 of water Use the enthalpy of combustion of methane in the data booklet to calculate the temperature rise of the water. 1. 0.25g of ethanol, C2H5OH, was burned and the heat given out raised the temperature of 500 cm3 of water from 20.1oC to 23.4oC. Calculate the enthalpy of combustion of ethanol -1269 kJ mol-1 2. 0.01 moles of methane was burned and the energy given out raised the temperature of 200cm3 of water from 18oC to 28.6oC. Calculate the enthalpy of combustion of methane. -1 -886.2 kJ mol 3. 0.1g of methanol, CH3OH, was burned and the heat given out used to raise the temperature of 100 cm3 of water at 21oC. Use the enthalpy of combustion of methanol in the data booklet to calculate the final temperature of the water. 26.4oC 4. 0.2g of methane, CH4, was burned and the heat given out used to raise the temperature of 250 cm3 of water Use the enthalpy of combustion of methane in the data booklet to calculate the temperature rise of the water. 10.66oC Tutorial Question – Self Check Page 10 of Tutorial booklet; 1. ΔH of butanol = -2664 kjmol -1 2. ΔH of ethanol = -1345 kjmol -1 3. ΔH of sulphur = -295 kjmol -1 4. ΔH of methane = - 877 kjmol -1 Now check your Combustion Chemical Equations! b) Calculation of the mass of products: Balanced Equation Calculations Learning intention Learn how the theoretical mass or volume of product can be calculated from the balanced reaction equation. From previous studies You should be able to • Write Formulae • Calculate percentage composition • Calculate the number of moles in a given mass • Calculate the number of moles of solute dissolved in a solution • Use balanced equations to work out unknown masses Reacting Masses Magnesium + oxygen….how much magnesium oxide is made? 1. Accurately weigh a crucible 2. Add approx 1.2g of Mg ribbon and reweigh 3. Place the crucible and lid in a silica triangle. 4. Heat gently at first then more strongly. Lift the lid with tongs from time to time to admit more oxygen, but not enough to let out the magnesium oxide. Reacting masses 1. When the reaction is complete, the magnesium will not glow more brightly when the lid is raised 2. Allow the crucible to cool 3. Reweigh the crucible Calculations Using the balanced equation calculate the mass of MgO you would expect to be formed? 2Mg + O2 → 2MgO Reacting masses - Questions 1. What mass of zinc sulphate will be produced on adding 6.0g zinc to excess sulphuric acid? 14.9g 2. What mass of sodium carbonate will react completely with 100cm3 of nitric acid concentration 1 mol l-1? 5.3g d) Excess Learning intention Learn how to calculate how much of a particular reactant is in excess from the balanced equation. Excess You can use the relative numbers of moles of substances, as shown in balanced equations, to calculate the amounts of reactants needed or the amounts of products produced. A limiting reactant is the substance that is fully used up and thereby limits the possible extent of the reaction. Other reactants are said to be in excess. Calculations involving excess A limiting reactant is the substance that is fully used up and thereby limits the reaction going further. Other reactants are said to be in excess. As soon as one of the reactant in a chemical reaction is used up the reaction stops. Any other reactant which is left over is said to be ‘in excess’. The reactant which is used up fully (limiting reactant) determines the mass of product formed – SO IS MOST IMPROTANT in terms of how much product is formed! You can use the relative numbers of moles of substances (mole ratios), as shown in balanced equations, to calculate the amounts of reactants needed or the amounts of products produced. Calculations involving excess Starter a) Which reactant is the limiting reactant when 10g of calcium carbonate reacts with 100cm3 of 1 mol l-1 hydrochloric acid? Equation: CaCO3 + 2HCl → CaCl2 + H2O + CO2 b) Now calculate the mass of carbon dioxide formed. CaCO3 is in excess. CO2 = O.22g Calculations involving excess Worked example 2. 1.2g of magnesium was added to 100cm3 2 mol l-1 hydrochloric acid. Calculate the reagent in excess and therefore the limiting reactant? Acid in excess. Graphs and Rates of Reaction e.g. Zn + 2HCl ZnCl2 + H2 Zn in excess 2mol l-1 HCl 20oC 2mol l-1 HCl 40oC Faster, but same amount of gas produced Vol H2 cm 3 1 mol l-1 HCl 20oC Half the gas produced time /s HCl is limiting reagent Calculations for you to try. 1. The graph below was obtained when 1.0g of powdered zinc was added to excess hydrochloric acid 1.0 mol l-1, copy the graph and sketch a line to show what you would expect if the reaction was repeated using a) 2.0 mol l-1 HCl and 1.0g Zn b) 1.0 mol l-1 HCl and 0.75g Zn Vol H2 cm 3 time /s 2. a) Calculate the limiting reactant when 3.27g of zinc is reacted with 100cm3 of 2.0 mol l-1 hydrochloric acid. b) What mass of hydrogen gas will be produced? 0.1 g Calculations involving excess Examples. For each of the following reactions calculate which reagent is in excess? a) 4.86g magnesium added to 250cm3 2 mol l-1 hydrochloric acid HCl b) 2.7g aluminium added to 200cm3 1 mol l-1 hydrochloric acid Al c) 2.43g magnesium added to 200cm3 1 mol l-1 sulphuric acid H2SO4 d) 3.27g zinc added to 100cm3 0.2 mol l-1 hydrochloric acid. Zn Calculations involving excess and molar gas volume An experiment was carried out to measure the concentration of hypochlotite ions (ClO-) in a sample of bleach. In this experiment the bleach sample is reacted with excess hydrogen peroxide. H2O2(aq) + ClO-(aq) H2O(l) + Cl-(aq) + O2(g) By measuring the volume of oxygen given off, the concentration of the bleach can be calculated. 80cm3 of oxygen was produced from 5.0cm3 of bleach. Calculate the concentration of the hypochlorite ions in the bleach (Take the molar gas volume to be 24 Litres mol-1). Calculations for you to try. 1. What mass of calcium oxide is formed when 0.4 g of calcium reacts with 0.05 mole of oxygen? 2Ca + O2 2CaO 0.56g 2. What mass of hydrogen is formed when 3.27g of zinc is reacted with 25cm3 of 2 mol l-1 hydrochloric acid? Zn + 2HCl ZnCl2 + H2 0.05 g Bonding in compounds Overview Learn how the elements can form bonds in compounds. Van der Waals forces Learning intention An introduction to the variety of intermolecular forces which exist between molecules. Relating physical properties to intermolecular forces Learning intention Learn how to explain differences in physical properties such as viscosity, melting point and boiling point in terms of differences in strength of intermolecular forces. The Chemical Bond Types of Chemical Bonding in Compounds Intramolecular (Within – i.e. between atoms) Intermolecular (between molecules) Van der Waals Forces Note only Covalent molecular compounds Ionic (compound only) Covalent bond(element or compound) New ! Polar covalent bond (compound only) contain these. 1. Hydrogen bonding 2. Permanent Dipole-Permanent Dipole interactions 3. London dispersion forces Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic 100 to 500 Covalent 100 to 500 Hydrogen 40 Dipole-Dipole 30 Londons forces 1 to 20 Ionic Compounds Ionic Compounds Ions - metals lose electrons and form positive ions - non-metals gain electrons to form negative ions - electrons are transferred from metals to non-metals transfer - + Na + Cl Na atom + Cl atom (2.8.1) (2.8.7) Na Cl Na+ ion + Cl- ion (2.8) (2.8.8) Covalent Molecular Compounds Covalent Bonding Sharing electrons • takes place between non-metal and non-metal • shared electrons count as part of the outer shell of both Atoms • shared electrons attract the nuclei of both atoms • this attraction is called the covalent bond Hydrogen chloride H H Cl (linear) Cl HCl Ammonia H H H N H H (pyrimidal) NH3 H N Water O H H (bent) H2O H O H Draw electron dot cross diagrams for the following molecules and structural formula 1. SCl2 2. CO2 3. CH4 X X S X X H Cl-S-Cl O=C=O H C H H Structure of Ionic Compounds Ionic Compounds The positive and negative ions are attracted (electrostatic bond )to each other. Ionic bond (electrostatic attraction) Na+ Cl- A giant lattice structure is formed. Each Na+ ion is surrounded by 6 Cl- ions. While each Cl- is surrounded by 6 Na+ ions. Ionic bonding is the electrostatic force of attractionbetween positively and negatively charged ions. This ionic network compound has many ionic bonds so ionic compounds have high m.p.s Ionic Compounds A giant lattice structure is formed when each Na+ ion is surrounded by 6 Cl- ions and each Cl- ion is surrounded by 6 Na+ ions. Sodium Chloride The formula of sodium chloride is NaCl, showing that the ratio of Na+ to Cl- ions is 1 to 1. The m.p. of NaCl is 801 0C The size of the ions will effect the strength of the ionic bond and how the ions pack together. e.g. NaF m.p. 1000oC, NaI 660oC Molecular Ions, e.g. SO4 Oxygen A single covalent bond. Sulphur O 2 additional electrons e.g. Copper can donate the extra 2 electrons needed. Cu Cu 2+ + 2e O S O O Copper sulphate contains the Cu2+ and the SO42- ions. There is, therefore, covalent bonding and ionic bonding in copper sulphate A solution of copper sulphate can conduct electricity. Molten ionic compounds can also conduct electricity. Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic 100 to 500 Covalent 100 to 500 Hydrogen 40 Dipole-Dipole 30 London’s forces 1 to 20 Discrete Covalent Molecular Compounds Covalent Molecular Compounds Discrete molecules are formed when two or more atoms share electrons. The atoms are non-metal elements. An example is methane. Methane: CH4 H H H C H H H C H H Methane has strong intra-molecular and weak inter-molecular. It’s b.p. is -183oC Covalent Molecular Compounds Non- metals elements can form double and triple covalent bonds. ethane C2H6 H H H C C H H H ethene C2H4 H H H C C H H H H H C C H H H H C C H H H C Double covalent bond Covalent molecular compounds have low m.p.’s because the weak forces holding the molecules together require only small amounts of thermal energy to break them. Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic 100 to 500 Covalent 100 to 500 Hydrogen 40 Dipole-Dipole 30 London’s Forces 1 to 20 Covalent Molecular Compounds Properties Low m.p.’s and b.p.’s., this increases with size of the molecule and the increasing number of atoms in the molecule. m.p.’s of the carbon halides 171 Temp / oC 90 -23 -183 CF4 CCl4 CBr4 CI4 m.p.’s increase because the strength of the London dispersion forces increase with the increasing size of the molecule. So more Energy is needed to separate molecules. Covalent Network Compounds Silicon Dioxide SiO2 Silicon and oxygen make up nearly 75% of the Earth’s crust. They are therefore the most common elements in the Earth’s crust. They combine together to make a covalent network compound called silicon dioxide. This is usually found in the form of sand or quartz. Each Si atom is bonded to 4 O atoms, and each O atom is bonded to 2 Si atoms. Hence the chemical formula, SiO2 . Silicon dioxide (silica) also has a high m.p. (1713 oC) and like SiC, it is very hard and used as an abrasive. It is relatively un-reactive. New Higher Chemistry E Allan J Harris Silicon Carbide SiC Silicon, like carbon, can form giant covalent networks. Silicon carbide exist in a similar structure to diamond. Tetrahedral shape C Covalent Bond C Si C C The 4 carbon atoms are available to bond with another 4 silicon atoms. This results in a COVALENT NETWORK COMPOUND Silicon Carbide SiC Silicon carbide (carborundum) has a chemical formula is SiC. As this compound is linked by strong covalent bonding, it has a high m.p. (2730oC). It is a hard substance as it is very difficult to break the covalent lattice. SiC is used as an abrasive for smoothing very hard materials. Each Si is bonded to 4 C’s and each C is bonded to 4 Si’s. Hence the chemical formula, SiC 3.32 Periodic trends in electronegativity Learning intention • Learn the definition of electronegativity • how to explain periodic trends in terms of nuclear charge, covalent radius and the screening effect of inner electrons. Electronegativity Electronegativity is a numerical measure of the relative ability of an atom in a molecule to attract the bonding electrons towards itself. Electronegativity Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond. e C e H Which atom would have a greater attraction for the electrons in this bond and why? Linus Pauling Linus American TodayPauling, we stillanmeasure chemist (and winner of of two electronegativities Nobel prizes!) up with the elements usingcame the Pauling concept scale. of electronegativity in 1932 to help explain the nature of chemical bonds. Since fluorine is the most electronegative element (has the greatest attraction for the Values for electronegativity bonding electrons) he assigned can be found on page 11 of it a value and compared all the data book other elements to fluorine. Electronegativities Looking across a row or down a group of the periodic table we can see a trend in values. We can explain these trends by applying the same reasoning used for ionisation energies. Looking across a period Increasing Electronegativity Li Be 1. 1.5 0 B C N O F 2. 2. 3. 3. 4. 0 5 0 5 0 What are the electronegativities of these elements? Across a period electronegativity increases The charge in the nucleus increases across a period. Greater number of protons = Greater attraction for bonding electrons Decreasing Electronegativity Looking down a group F 4.0 Cl 3.0 What are the electronegativities of these halogens? Br 2.8 I 2.6 Down a group electronegativity decreases Atoms have a bigger radius (more electron shells) The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells. Trends in electronegativity Electronegativity increases across a period. Electronegativity decreases down a group Going across the period, the nuclear charge increases. This pulls the electron shells closer to the nucleus. As a results, the electronegativity increases. Going down the group, the nuclear charge increases but the number of electron shells also increases. As a result of ‘shielding’ and an increase distance the outer shell is from the nucleus, electronegativity decreases. Chemical bonds: types of bonds Explores how different types of bonds are formed due to variations in the electronegativity of the bonded atoms. The distortion of the orbitals and the polarity of the bond is also displayed. Linus Pauling (1901-1994) An account of the life and work of the Nobel Prize-winning chemist, Linus Pauling. Periodic Table of Data Visual database of the physical and thermochemical properties of the chemical elements which allows the user to plot graphs and tables, play games and view diagrams. The Bonding Continuum Predicting Non-Polar, Polar and Ionic Bonds Learning intention Learn how differing differences in electronegativity between bonding atoms lead to the formation of non-polar covalent bonds, polar covalent bonds and even ionic bonds. Bonding continuum Learning intention Learn about the bonding continuum which stretches between pure covalent and ionic bonds, in terms of differences of electronegativity between bonding atoms. Covalent Bonding A covalent bond is a shared pair of electrons Both nuclei try to pull the electrons towards electrostatically attracted to the positive nuclei themselves of two atoms. + - + The achieve a stable This is atoms like a tug-of-war whereouter bothelectron sides are arrangement gas arrangement) by pulling(aonnoble the same object. sharing It creates a strong bondelectrons. between the two atoms. Covalent Bonding Picture a tug-of-war: If both teams pull with the same force the midpoint of the rope will not move. Pure/Non-Polar Covalent Bonds No/Very small difference on Electronegativity H e e H This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms. All non-metal elements have Pure/Non-polar covalent bonds. Polar Covalent Bonds Non-polar covalent bond – electrons shared equally between atoms (same electronegativity) Polar covalent bond – electrons shared unequally between atoms (atom B is more electronegative) Polar Covalent Covalent Bonds Bigger difference in electronegativity What if it was an uneven tug-of-war? The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right. Polar Covalent Bond A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally. This is due to different elements having different electronegativities. Polar Covalent Bond δ- e.g. Hydrogen Iodide δ+ H e e I If hydrogen iodide contained a pure covalent bond,makes the electrons would negative be shared equally as This iodine slightly and hydrogen shown above. slightly positive. This is known as a dipole. However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war) Polar Covalent Bond In general, the electrons in a covalent bond are not equally shared. e.g. C δ+ 2.5 Cl δ- 3.0 Electronegativities δ- indicates where the bonding electrons are most likely to be found. Polar Covalent Bond Consider the polarities of the following bonds: Difference Electronegativities Bond C Cl 2.5 3.0 0.5 P H 2.2 2.2 0 O H 3.5 2.2 1.3 P H δ+ C δ- Cl δ- O δ+ H Increasing Polarity Complete a similar table for C-N, C-O and P-F bonds. Polar Covalent Bonds In the covalent bond between fluorine and hydrogen. The bonding electrons are not shared equally between the two atoms. Fluorine Hydrogen The fluorine nucleus has more protons and has a stronger pull on the electrons than the hydrogen nucleus.. d- F 4.0 d+ H 2.2 Thus the fluorine atom has a greater share of the bonding electrons and acquires a slight negative charge. The hydrogen atom is then made slightly positive. The bond is a polar covalent bond and we use the symbols d+ and d- to show this. The dipole produced is permanent. Fluorine is the most electronegative element. It is small atom compared to others and its nucleus is massive for its atomic size. Some other polar covalent bonds are O-H and N-H. They have large differences in Electronegativty. Bonding Continuum “Covalent compounds are formed by non-metals only” IS NOT AN ABSOLUTE LAW! Some compounds break this rule…. The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character. Electronegativity Difference and Bond Type: Difference Bond Example 0.0-0.5 Non-Polar/(pure)Covalent Bond H-H 0.0 0.5-1.5 Polar Covalent Bond H-Cl H20 0.9 0.7 > 1.5 Ionic NaCl 2.1 *Note Exceptions: Very Polar bonds 1.5-2.0 Very Polar Covalent (almost ionic) H-F 1.9 Making Tin(IV)iodide • Gently heat the tin and iodine in a small conical flask containing 10cm3 of tolulene on a hot plate. • Collect the yellow precipitate by filtration using Büchner filtration Making Tin(IV)iodide • Determine the melting point of the solid collected. Making Tin(IV)iodide Melting point of tin(IV)iodide is 143oC. Tin electronegativity of 1.8 Iodine has electronegativity of 2.6 Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therefore only weak London dispersion forces so low melting an boiling point. Titanium (IV) chloride TiCl4 is a dense, colourless liquid. It is one of the rare transition metal halides that is a liquid at room temperature, This property reflects the fact that TiCl4 is ………….; that is, each TiCl4 ………. is relatively …………… associated with its neighbours. Most metal chlorides are ionic. The attraction between the individual TiCl4 molecules is weak, primarily ……………….……….., and ……………. these weak Van der Waals (intermolecular) interactions result in low melting and boiling points. TiCl4 is soluble in toluene and dichloromethane, as are other non-polar species. TiCl4 • Used in smoke grenades and for smoke screens Non-Polar/Polar Bonds vs Non-Polar/Polar Molecules • We have just learned about polar and nonpolar bonds and how to identify them. • But just because a molecule has polar bonds, doesn’t mean it is overall a polar molecule. • Checks have top be done! Symmetry CCl4 Cl d Symmetrical molecule d+ Cl Cl d d Cl Tetrachloromethane has a symmetrical d arrangement of polar 4 polar covalent C-Cl bonds and the polarity bonds in CCl4 tetrahedral cancels out over the shape molecule. NON-POLAR molecule Symmetry CO2 Symmetrical molecule d- O d+ O d- 2 polar covalent C=O bonds in CO2 linear shape NON-POLAR molecule Carbon dioxide has a symmetrical arrangement of polar bonds and the polarity cancels out over the molecule. Symmetry CHCl3 Check 1: It has polar bonds Check 2: It is not symmetrical H Permanent dipole Asymmetrical molecule d+ Cl d- Cl Cl d- d- 3 polar covalent C–Cl bonds and 1 polar covalent C-H bond in CHCl3 It is overall a POLAR molecule London Dispersion Forces These Only exist in Non-Polar Molecules. Remember; London Dispersion forces are weak forces of attraction between molecules. •They only exist between non-polar molecules. •All non-metal discrete covalent molecular molecules are non-polar •We now know that many covalent molecular compounds are non-polar •Non Polar molecules have low m.p. And b.p due to the weak London forces between the molecules. A dipole can induce other atoms to form dipoles, resulting in temporary dipole –dipole attraction. Londons forces Permanent dipole-permanent dipole interactions Learning intention Learn about this additional intermolecular force of attraction which exists between polar molecules. Permenant Dipole-Dipole Attractions – Only in Polar Molecules The differing electronegativities of different atoms in a molecule and the spatial arrangement of polar covalent bonds can cause a molecule to form a permanent dipole – i.e. It is a polar molecule.. H Permanent dipole Asymmetrical molecule d+ Cl d- Cl Cl d- d- 3 polar covalent C–Cl bonds and 1 polar covalent C-H bond in CHCl3 POLAR molecule Permanent Dipole-Dipole interactions Molecules with permanent dipoles attract each other. The attraction is stronger than Londons forces Note: Hydrogen bonding is a particular example of dipole-dipole attractions –we will see this later. Comparing Properties of Polar and Non Polar Molecules Both propanone and butane have the same formula mass of 58 however, butane boils at – 1 oC while propanone boils at 56oC Propanone is a polar molecule as it has a permanent dipole, so has polar-polar attraction as well as London’s forces between molecules. H H C H d- O H C C d+ H H b.p. 56 o C Butane has no permanent dipoles, so only London’s forces between molecules. So has a lower boiling point. H H C H H C H H C H H C H H b.p. -1 o C Hydrogen Bonding Learning intention Learn about this strong type of intermolecular forces which exists between molecules containing N-H, O-H or F-H bonds. Relating physical properties to intermolecular forces Learning intention Learn how to explain differences in physical properties such as viscosity, melting point and boiling point in terms of differences in strength of intermolecular forces. Intermolecular - Hydrogen Bonding Consider the compounds formed between elements in group 4 of the Periodic table and hydrogen The group 4 hydrides are CH4, SiH4, GeH4, SnH4 They are all covalent molecular so have low melting points and boiling points. Boiling Point (K) Group 4 250 200 150 100 50 0 Group 4 CH4 SiH4 GeH4 SnH4 The boiling point increases as you go down the group. As you go down the group the central atom gets bigger. There are more electrons so a greater chance of an uneven distribution of electrons within the atom. The London’s forces between the molecules gets stronger as you go down the group. More energy is needed to separate the molecules from each other. Intermolecular – Hydrogen Bonding A similar pattern would be expected in the other covalent molecular hydrides The group 5 hydrides NH3, PH3, AsH3 and SbH3 The group 6 hydrides H2O, H2S, H2Se and H2Te The group 7 hydrides HF, HCl, HBr and HI Boiling Point (K) Group 5 300 200 Group 5 100 0 NH3 PH3 AsH3 SbH3 NH3, has a higher boiling point than expected. Boiling Point (K) Group 6 400 300 200 Group 6 100 0 H2O H2S H2Se H2Te H2O has a higher boiling point than expected. Boiling Point (K) Group 7 400 300 200 Group 7 100 0 HF HCl HBr HI HF has a higher boiling point than expected. Intermolecular - Hydrogen Bonding Boiling Points of Hydrides Boiling Point (K) 400 H O 2 300 HF Group 4 Group 5 Group 6 Group 7 NH3 200 100 0 Series Number It is more difficult to separate NH3, H2O and HF molecules from each other than expected. Intermolecular - Hydrogen Bonding These compounds all have H atoms directly bonded to very electronegative atoms. But only if electronegativity beween H and other atom is big enough, to produce a large dipole, will Hydrogen bonding be possible. In HF the H-F bond is polar covalent. The F has a much higher electronegativity than H. The pair of shared electrons in the covalent bond spend more time closer to the fluorine than the hydrogen. The H-F bond is very polar. Hδ+ - Fδ- Intermolecular - Hydrogen Bonding The HF molecules can attract each other Hδ+ - Fδ- Hδ+ - Fδ- Hδ+ - Fδ- This is called hydrogen bonding. Hydrogen bonding is weak but is stronger than very weak London forces. Intermolecular - Hydrogen Bonding NH3 has H atoms directly bonded to very electronegative N atoms. Nd- Hd+ Hd+ Hd+ Hd+ Nd- Nd- Hd+ Hd+ Hd+ Hd+ Hd+ There are Hydrogen bonds as well as London forces between the ammonia molecules. Intermolecular - Hydrogen Bonding H2O has H atoms directly bonded to very electronegative O atoms. Od- Hd+ Hd+ Od- Hd+ Od- Hd+ Hd+ Hd+ There are Hydrogen bonds as well as London forces between the water molecules. Proteins consist of long chain atoms containing polar C=O and H-N bonds. Hydrogen bonds help give enzymes their shape. Water d- O H H d+ d+ Oxygen has 2 lone pairs of electrons which can form a hydrogen bonds with two hydrogen atoms. Each water molecule, in theory, could be surrounded by 4 hydrogen bonds. Hydrogen bonding in water Density of water Water Water has its greatest density at a temperature of 4oC. When, as water cools further, the molecules start to move further apart, due to the hydrogen bonding, until a more open structure is formed at its freezing point. So ice floats!! New Higher Chemistry E Allan J Harris Hydrogen bonding in ice Hydrogen bonding in solid water gives rise to an open structure. This is why ice is less dense than liquid water. Hydrogen bonding is also responsible for holding the two strands of nucleic acids together in DNA Viscosity Density of water Viscosity is related to the molecular mass and the number of – OH present. Hydrogen bonding between the molecules will increase its viscosity. New Higher Chemistry E Allan J Harris Water Surface tension Water has a high surface tension. The molecules on the surface have in effect, hydrogen bonds. This has the effect of pulling the surface molecules closer together. Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic 100 to 500 Covalent 100 to 500 Hydrogen 40 Dipole-Dipole 30 London’s forces 1 to 20 Behaviour in electrical fields New Higher Chemistry E Allan J Harris Nappies • Cloth nappies cost between £100-£400 as opposed to disposable at £800-£1,200 for the 2.5 years of normal nappy use. • 3 billion nappies are thrown away in the UK each year with 90% going to landfill. They can take up to 500 years to decompose. • Disposables make up 4% of total household waste and up to 50% of that of families with one baby • Disposable nappies use up to 5 times more energy to produce than cotton ones – that's including the washing process . • Seven million trees are felled every year in Canada and Scandinavia to supply the pulp for disposables sold in the UK. Sodium polyacrylate is a polymer with a molecular weight of over one million! sodium carboxylate Chemical Background Groups called sodium carboxylate are attached along the backbone. Sodium poly(acrylate) absorbs 500 times its own mass of water. water + + Na+ - Sodium poly(acrylate) absorbs 500 times its own mass of water. + + ++ Sodium poly(acrylate) absorbs 500 times its own mass of water. - - - - - Predicting solubility from solute and solvent polarities Learning intention Learn how the polarity of both the solute and solvent molecules influences solubility. Solvent Action Solvent Action A liquid that a substance dissolves in is called a SOLVENT. Solvents can be either polar or non-polar molecules. Immiscible liquids do not mix, e.g. oil and water, however, non-polar liquids are miscible with each other. Polar solvents will usually dissolve polar molecules. Non-polar solvents will usually dissolve non-polar molecules. Water is a polar molecule so it is a polar solvent. d+ H Water has a polar covalent bonding between O and H. H d+ d+ O d- dd+ Dissolving in Water Ionic Compound dissolving in water: d- d+ d+ - d+ + - d+ + - d+ d- + ddd+ d+ + d+ d- d+ d+ Hydrated ions dd+ d+ d- d+ - d+ d+ d+ d- Dissolving in Water Dissolving in Water Pure Hydrogen chloride is polar covalent. When water is added it breaks to produce ions dd+ H d- Cl d+ d+ dd+ d+ d+ d+ d- H+ ddd+ d- d+ d+ d+ Hydrated ions d+ d+ d- d+ Cl- d+ d+ d+ d- Dissolving in Water Generally, covalent molecules are insoluble in water. However, small moleculeslike ethanol (C2H5OH), with a polar O-H functional group, will dissolve, dd+ H2O d+ Ethanol H H H C C H H d- O H d+ Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic 100 to 500 Covalent 100 to 500 Hydrogen 40 Dipole-Dipole 30 London’s forces 1 to 20