Jimmy's Combo Chemistry PDF

Check list Dr. Muhammed Gamal Topic 1 Yes Avg. No 1. Definition of Avogadro’s Constant. 2. Definition of Mole. 3. How to calculate: a. N...

Check list Dr. Muhammed Gamal Topic 1 Yes Avg. No 1. Definition of Avogadro’s Constant. 2. Definition of Mole. 3. How to calculate: a. Number of Molecules b. Number of Atoms c. Number of Ions d. Number of Protons, Neutrons and Electrons 4. When and How to use mole rules: a. n=m/Mr b. n=CV c. n=V/24 5. When and How to use the Mole Ratio. 6. Using m/Mr=m/Mr, when No. of Moles are equal. 7. How to use Molar Ratio to indicate which substance is in Excess. 8. How to use Molar Ratio to get volume of gases. 9. Know the Difference between the Volume of gases produced and gases remaining. 10. How to calculate Percentage Yield and Percentage Purity. 11. How to calculate Atom Economy. 12. How to get the Empirical Formula. 13. How to get the Molecular Formula from Empirical Formula. 14. How to obtain Water of Crystallisation in hydrated Crystals: a. by Molar Ratio Rule. b. by calculating the Mr of the whole hydrated crystals. 1  Check list Dr. Muhammed Gamal 15. What is the ideal gas rule and what are the measurements’ SI units. 16. How to convert between ℃ and K. 17. What is the Difference between the Molar Concentration and Mass Concentration. 18. What is the Definition of Parts per million (PPM). 19. What is the PPM rule. 2 Unit 1 Topic 2 …………..Summary Toqa Khodair  Definitions Atomic number: the number of protons in the nucleus of an atom. Mass number: the sum of the number of protons and the number of neutrons in the nucleus of an atom. Isotopes: atoms of the same element that have the same atomic number but different mass number. Relative atomic mass: the weighted mean mass of an atom of the element compared to 1/12 of the mass of an atom of carbon-12. Relative isotopic mass: the mass of an individual atom of a particular isotope relative to 1/12 of the mass of an atom of carbon-12. Molecular ion peak: the peak with the highest m/a ratio in the mass spectrum, the M peak. Quantum shell: the energy level of an electron. Orbital: a region within a atom that can hold up to two electrons with opposite spins. Electronic con guration: the number of electrons in each sub-shell in each energy level of the atom. Hund’s rule: electrons will occupy the orbitals singly before pairing takes place. Pauli exclusion principle: two electrons cannot occupy the same orbital unless they have opposite spins; electrons spin is usually shown by using upward and downward arrows. First ionisation energy: the energy required to remove an electron from each atom in one mole of atoms in the gaseous state. Second ionisation energy: the energy required to remove an electron from each singly charged positive ion in one mole of positive ions in the gaseous state. Third ionisation energy: the energy required to remove an electron from each doubly charged ion in one mole of positive ions in the gaseous state. Groups: the vertical columns in the periodic table. Periods: the horizontal rows in the periodic table. Periodic properties: regularly repeating patterns of atomic, physical and chemical properties, which can be predicted using the periodic table and explained using the electron con gurations of the elements. fi fi  Top tips S-orbitals have a spherical shape while p-orbitals have dumbbell shape. S-block elements are group 1 and 2 elements. P-block elements are group 3, 4, 5, 6, 7, and 8 elements. D-block elements are the transition metals. When using the s, p, d notation, the sub shells are lled in order. S, then p then d. Copper and chromium are the exceptions to the electronic con guration. 10 Cu: 1s2 2s2 2p6 3s2 3p6 3d 4s1 5 Cr: 1s2 2s2 2p6 3s2 3p6 3d 4s1 Successive ionisation energies increases steadily and big jumps occur at de ned places. The big jumps occur due to the electrons being removed from shells of lower energy level. This is because as the ion gets smaller, there is greater effect of nuclear charge, the ionic radius gets smaller, distance and shielding decreases. Ionisation energy decreases down the group as the radius increases (distance and shielding increases), the increase in nuclear charge however is less signi cant. Across the period, ionization energy generally increases as the nuclear charge increases, shielding remains the same, distance decreases, and so there is greater attraction between the outermost electron and the nucleus. Group 3 and group 6 elements are anomalies. Group 3 elements have their outermost electron in a p sub shell, which is in a higher energy level compared to electrons in s-sub shell. So, less energy is needed to remove the electron. For example: beryllium (group 2) and boron (group 3). Boron has greater nuclear charge compared to beryllium however, requires less energy has it has its outermost electron in p orbital, which has higher energy. Group 6 elements have paired electrons in 2p orbital thus increasing the electron-electron repulsion and so less energy is needed to remove the electrons. For example: nitrogen (group 5) and oxygen (group 6). Oxygen has 2 electrons in a p orbital, which increases the electron-electron repulsion. So, less energy is required to remove the electron. Atomic radius decreases across the period as nuclear charge increases so greater attraction between nucleus and outermost electrons. It increases down a group as more shells are being added. Atomic radius across periods First Ionisation energy across a period fi fi fi fi  The mass spectrometer is used to determine masses of compounds, carbon/hydrogen dating, use of illegal drugs by athletes, and pharmaceutical industry. There are ve steps to be done when using mass spectrometry- vaporization, ionisation, acceleration, de ection, and nally detection. For the sample to be able to move through the mass spectrometer, it needs to be in gaseous state. The vapor is then bombarded with high energy electrons using an electron gun to form positive ions to be detected by the detector. It is then accelerated using negatively charged particles, and de ected according to the mass to charge ratio as they pass through a uniform magnetic eld. Greater mass to charge ratio will be de ected less than smaller mass to charge ratio. The ions are then detected by a detector. The ratio of chlorine 35 to chlorine 37 is 3:1 and the ratio of chlorine molecules (70, 72, 74) is 9:6:1. Mass spectrometer fl fi fi fl fi fl  Check list Dr. Muhammed Gamal Topic 2 Yes Avg. No 1. What is the Mass and Charge of each Sub-Atomic Particle. 2. Understanding the Structure of the Atom: a. Quantum Shells b. Sub-Shells c. Orbitals 3. Understanding the trend in the Energy between Sub-shells. 4. Understanding the trend in the Energy between Quantum Shells. 5. Understanding the No. and which Sub-shells are found in each Quantum Shell. 6. Understanding and Memorising the Shape of S and P orbitals. 7. How many electrons are held by each orbital and how they are held. 8. How many orbitals are found in each Sub-Shell. 9. How many electrons are found in each Sub-Shell. 10. What is the order that the electrons are arranged in to fill up the Sub-Shells. 11. Understanding how to perform the electronic configuration of elements. 12. Understanding that the electrons are removed first from the Sub-Shell with the highest energy. 13. Understanding the s,p,d and f blocks positions in the periodic table. 14. Understanding why and which block is meant for each group. 15. Understanding how to write the electronic configuration of 24Cr. 16. Understanding how to write the electronic configuration of 29Cu. 17. The Definition of the ionisation energy. 1  Check list Dr. Muhammed Gamal 18. Understanding How to write the equation of the First, Second and Third ionisation energies of an element. 19. What is the Trend of the first ionisation energy as we go down in Group 1. 20. What is the relation between reactivity and ionisation energy. 21. What is the trend of the ionisation energy as we go down the group. 22. What is the trend of the ionisation energy as we go right across the period. 23. What does the big Jumps in the ionisation energies indicate. 24. Understanding why does the Successive ionisation energies increase: a. Nuclear Charge (No. of Protons). b. Attraction Force (Distance between Electrons and Nucleus). c. Shielding (No. of energy Shells). d. Atomic/Ionic Size (Nuclear Charge and Shielding). e. Electron-Electron Repulsion. 25. Understanding why First Ionisation energy decreases as we go down Group 1. 26. Understanding why First Ionisation energy increases as we go right across the period. 27. Understanding Why Group 3 doesn’t follow the general increase in the ionisation energy. 28. Understanding Why Group 6 doesn’t follow the general increase in the ionisation energy. 29. Understanding the Trend of Atomic Radius across the Period. 30. Understanding the Trend of Atomic Radius across the Group. 31. Understanding the Trend of Melting Point across the Period. 32. Understanding the Trend of Melting Point across the Group. 33. Understanding the Definition of Isoelectronic. 2  Check list Dr. Muhammed Gamal 34. Understanding what we can use Mass Spectrometer in. 35. What is the Mass Spectrometer’s Applications. 36. Understanding How does Mass Spectrometry work: a. Vaporisation b. Ionisation c. Acceleration d. Deflection e. Detection 37. Understanding Why does Acceleration is important in Mass Spectrometer. 38. Understanding How to write the Equation of Ionisation phase in Mass Spectrometry. 39. What Element can be used to detect the leaking in Mass Spectrometer, and why. 40. Understanding the Purpose of the Vacuum Pump in Mass Spectrometer. 41. Understanding How to calculate Relative Atomic Mass. 42. Understanding the Ratio between Chlorine atoms (35Cl & 37Cl). 43. Understanding the Ratio between Chlorine Molecules (Cl2) and How to calculate this Ratio. 44. Understanding the Ratio between Bromine atoms (79Br & 81Br). 45. Understanding the Ratio between Bromine molecules (Br2). 46. Understanding that molecule/Atom could have a +2 Charge in Mass Spectrometry. 47. Understanding How to switch between log values and Real values. 48. Understanding that only Positively Charged Species are detected. 49. Understanding What does Molecular ion peak refer for. 3 Unit 1 Topic 3 …………. Summary Toqa Khodair  Glossary Ionic bonding: the electrostatic forces of attraction between oppositely charged ions arranged in a regular lattice. Polarizing power: the ability of a positive ion to distort the electron density of a neighboring negative ion. Polarization: the distortion of the electron density of a negative ion. Electronegativity: the ability of an atom to attract a bonding pair of electrons in a covalent bond. Polar covalent bond: a type of covalent bond between two atoms where the bonding pair of electrons are unequally distributed; because of this, one atom carries a slight negative charge and the other a slight positive charge. Discreet molecule: an electrically neutral group of two or more atoms held together by covalent bonds. Dative covalent bond: the bond formed when an empty orbital of an atom overlaps with an orbital containing a lone pair of electrons of another atom. Electron pair repulsion theory: the electron pairs on the central atom of a molecule or ion arrange themselves in order to create the minimum repulsion between them; lone pair – lone pair repulsion is greater than lone pair-bond repulsion which in turn is greater than bond pair-bond pair repulsion. Dipole: exists when two charges of equal magnitude but opposite signs are separated by a small distance. Delocalized electrons: electrons that are not associated with any single atom or any single covalent bond. Metallic bonding: the electrostatic force of attraction between the metal cations and the delocalized electrons. Ionic bonding Ionic bonding occurs in solid materials consisting of a regular array of oppositely charged ions extending through a giant lattice network in other words it’s the electrostatic attraction between oppositely charged ions arranged in a regular lattice. Dot and cross diagrams Sodium chloride Magnesium oxide + - 2 2 Na Cl Mg O Evidence for the existence of ions When passing direct current through copper(II) chromate (VI) solution. At the negative electrode blue color of copper ions will be seen and at the positive electrodes yellow color will be seen. This is because copper ions migrate towards the negative electrode and the solution around this electrode turns blue while the chromate ions migrate towards the positive electrodes and the solution around this electrode turns yellow. Another piece of evidence is the electron density map. The electron density map of ionic compounds clearly show a gap which separates the ions. Polarization Polarization is the distortion of the electron density of a negative ion. The polarizing power is the ability of a positive ion to distort the electron density of a neighboring negative ion. Polarization will be increased by: Cation Anion Higher charge Higher charge Smaller size Larger size  Properties of ionic compounds 1) high melting point 2) brittleness 3) poor electrical conductivity when solid 4) good electrical conductivity when molten 5) soluble in water Solid ionic compounds do not conduct electricity as there are no delocalized electrons or ions that are able to move under the in uence of an applied potential difference. However in aqueous solution or molten state the ionic compounds will conduct electricity since the ions are now mobile and will migrate to the electrodes of opposite sign when a potential difference is applied. Covalent bond Covalent bond forms between two atoms atomic orbital containing a single electron from one atom overlaps with an atomic orbital, which also contains a single electron, of another atom. There are three ways an orbital can overlap: An end on overlap leads to the formation of sigma bonds. A sideways overlap of p orbitals leads to the formation of pi bonds. Pi bonds cannot form until a sigma bond has been formed. For this reason pi bonds only exist between atoms that are joined by double or triple bonds. Pi vs sigma bonds Pi bonds are weaker than sigma bonds. Sigma bonds are found along a straight line across the nucleus of both atoms but pi bonds are parallel and are found in two places. Pi bonds are caused by a sideways overlap of 2p orbitals. Bond length and bond strength Bond length is the distance between the nuclei of two atoms that are covalently bonded. Bond strength is the amount of energy required to break one mole of bond in gaseous state. As bond length decreases, bond strength increases as the electrostatic forces between the two nuclei and the electrons in the overlapping atomic orbitals. Electronegativity The ability of an atom to attract a bonding pair of electrons. fl Trend In electronegativity Across period The electronegativity increases across a period. Down a group Electronegativity decreases down a group. Electron density map Dot and cross diagram Dative covalent bonds A dative covalent bond is formed when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair of electrons of another atom. The bond is often represented by an arrow from the atom providing the pair of electrons to the atom with a empty orbital. Examples: The hydroxonium ion The ammonium ion Aluminum chloride Shapes of molecules Shape and polarity Non-polar; dipoles cancel out Non-polar; dipoles cancel out Non-polar; dipoles cancel out Polar; dipoles reinforce one another Polar; dipoles reinforce one another Polar; dipoles reinforce one another Properties of metals Metallic bonding High melting temperatures The electrostatic forces of attraction between the cations and the delocalized electrons. Good electrical conductivity Good thermal conductivity Malleability Ductility Reasons for these properties Metals have a giant lattice structure and so needs lots of energy to overcome the forces of attraction between the cations and the delocalized electrons. The number of the delocalized electrons per cation plays in important part in determining the melting temperature. Metals are a good electrical conductor as they have delocalized electrons that are free to move and carry current. Metals can be hammered or pressed into different shapes. They can also be drawn into a wire. Both of these properties depends on the ability of the delocalized electrons and the cations to move throughout the structure. When stress is applied to a metal, the layers of the cations may slide over one another however because the delocalized electrons are free moving, they move with the cations and prevents strong forces of repulsion forming between the cations in one layer and the cations in a different layer. Giant covalent lattice Diamond Diamonds are extremely hard because of the very strong carbon to carbon bonds in throughout the structure. It has a very high melting temperature because of the great number of strong carbon to carbon bonds that have to be broken in order to melt it. Graphite Graphite is a good conductor of electricity as it has delocalized electrons between the layers which can carry currents and move freely. Graphite has a high melting temperature for the same reason as diamonds. Graphene Graphene is a pure carbon in the form of a very thin sheet, one atom thick. It’s also can be described as a one atom thick layer of graphite. Uses of Graphene: Properties of Graphene Flexible electronics Thin Microelectronics Flexible Batteries Transparent Making nanotubes Oxidation resistant Solar panels Low density Touchscreens Durable High melting point Summary  Check list Dr. Muhammed Gamal Topic 3 Yes Avg. No 1. Understanding the Definition of Ionic Bonding. 2. Understanding How Ionic Bonding are established. 3. Understanding the Factors that affect the strength of the Ionic Bonding: a. Size of Ions (Anions & Cations). b. Charge of Ions (Anions & Cations). 4. Knowing how to provide Evidence for Ions Existence: a. Electrolysis b. Electron Density Map 5. Understanding the Increase in the Ionic Radii down the Group. 6. Understanding the Decrease in the Ionic Radii across the Period to the Right. 7. Understanding the Meaning of Polarisation and Polarising Power. 8. Understanding the Factors that affect the Polarisation: a. Size of Cations and Anions. b. Charge of Cations and Anions. 9. Understanding How an Ionic Compound could have a Covalent Character related to the Polarisation. 10. Understanding the Ionic Compound Properties: a. Melting and Boiling Points b. Water Solubility c. Electrical Conductivity d. Brittlely 11. Understanding the Definition of Covalent Bonding. 12. Understanding How Covalent Bonding are established due to Orbitals’ Overlapping. 1  Check list Dr. Muhammed Gamal 13. Understanding the Difference between Sigma and Pi Bonds. a. Way of Overlapping. b. Strength of Overlapping. 14. Knowing the No. of Sigma Bond and Pi Bonds in: a. Single Bond b. Double Bond c. Triple Bond 15. Understanding the Factors that affect the strength of the Covalent Bonding: a. Bond Length b. Bond Strength 16. Understanding the Meaning of Electronegativity. 17. Understanding the Factor that affect the Electronegativity. 18. Understanding the Decrease of Electronegativity down the group. 19. Understanding the Increase of Electronegativity across the period to the right. 20. Understanding How a Covalent Compound could have an Ionic Character related to the Electronegativity. 21. Understanding that Some Molecules/Compounds could have Zero net electronegativity. 22. Understanding the Definition of Dative Covalent Bond. 23. Understanding How Dative Covalent Bonds are established due to Orbitals’ Overlapping. 24. Understanding How differentiate between Dative Covalent bonds and Covalent Bond. 25. Knowing the Examples on Dative Covalent Bonds found in the Notes: a. Ammonium ion [NH4]+ b. Hydronium ion [H3O]+ c. Aluminium Chloride [Al2Cl6] 2  Check list Dr. Muhammed Gamal 26. Understanding How the molecules have different Shapes. 27. Understanding the Difference in Repulsion force between: a. Two Bonded Pairs. b. One Bonded Pair and One Lone Pair. c. Two Lone Pairs. 28. Knowing that Double Bond can be described as Bonded Pairs of Electrons. 29. Memorising the Angles and Type of each Shape associated with the No. of Bonded Pairs and Lone Pairs in each Shape: a. Linear b. Trigonal Planar c. Tetrahedral d. Trigonal Pyramidal e. V-Shaped f. Trigonal bipyramidal g. Octahedral 30. Understanding the Difference between the Polar and Non-polar Compounds. 31. Understanding the Meaning of Dipole Bonds. 32. Understanding that a compound could be Non-Polar with Polar Bonds. 33. Understanding the Definition of Metallic Bonding. 34. Memorising the Physical Properties of Metallic Bonding. 35. Understanding How Metals can conduct Electricity at any Phase. 36. Understanding Why Metals are Malleable and Ductile. 37. Understanding Why Metallic Bonding Increases as we go right across the period. 38. Understanding the Structure of Diamond (Same as Silicon). 39. Understanding Why Diamond/Silicon has High Melting and Boiling Points. 3  Check list Dr. Muhammed Gamal 40. Understanding the Structure of Graphite. 41. Understanding Why Graphite is Soft. 42. Understanding Why Graphite can conduct Electricity. 43. Understanding What is Graphene. 44. Memorising the Uses of Graphene. 4 Unit 1 Topic 4 …………..Summary Menna Mohammed Badawy  Definitions Hydrocarbon: A compound made up of hydrogen and carbon only. Saturated: A compound containing only single bonds. Unsaturated: A compound containing one or more double bonds. Displayed formula: shows every atom and every bond. Structural formula: Shows how the atoms are joined together. Skeletal formula: shows all the bonds between carbon atoms. Molecular formula: shows the count of each atom in a molecule. Empirical formula: the simplest whole number ratio of the molecular formula. Functional group: an atom or group of atoms in a molecule that’s responsible for its chemical reactions. Homologous series: a family of compounds with the same functional group, and differ in formula by CH2 for the next member. Structural isomers: Compounds with the same molecular formula but different structures. Homolytic fission: the shared pair of electrons in the covalent bond is divided equally between both atoms. A dot is used to represent the electron. Heterolytic fission: the breaking of a covalent bond so that both bonding electrons are taken by one atom. A positive ion and a negative ion are formed in the process. Electrophiles: a species attracted to a region of high electron density, normally positively charged. Free radicals: a species that contains an unpaired electron. Hazard: something that could cause harm to a user. Risk: the chance of a hazard causing harm. Fractional distillation: A mixture of liquids are separated according to their boiling points. Cracking: the breakdown of long-chained molecules into short-chained ones. Reforming: conversion of a straight chain hydrocarbon into a branched or a cyclic structure to allow them to burn more efficiently  Top tips Types of formula: Naming the compounds: When naming compounds, we use the smaller Locant values. If there is a functional group, we give the functional group the priority in naming the compound (Gets smaller locant value) Prefixes: Types of isomers: Chain isomers: are when the molecules have different carbon chains i.e. Butane and methylpropane. Position isomers: molecules with the same functional group but attached in different positions on the same carbon chain. Types of reactions: Addition reactions: When 2 reactant species combine to form a single product. Substitution reactions: When 2 reactant species react, and an atom/ion Substitutes another. i.e. C2H5Br + OH- —> C2H5OH + Br- Oxidation reactions: an organic compound oxidized by another reagent. The oxidized compound might gain oxygen or lose hydrogen. Reduction reactions: one organic compound is reduced, sometimes by hydrogen gas and a catalyst or an inorganic reagent. Organic compound can either gain hydrogen or lose oxygen. Addition Polymerization reactions: 2 or more reactant molecules are joined together to form one large molecule. Hazards and risks: Combustion of organic compounds: Incomplete combustion: Produces Carbon monoxide and water, sometimes produces carbon (soot). Complete Combustion: produces carbon dioxide and water. Oxides of Sulfur: Some crude oils contain traces of sulfur, these atoms then turn into sulfur dioxide. SO2 dissolves in water forming sulfuric acid which contribute to acid rain. Oxides of Nitrogen: high temperature and pressure in the engine cause the nitrogen in the air to react with oxygen. These products contribute to acid rain as well. Carbon neutrality: Amount of Carbon absorbed by the fuel source is equal to the amount of carbon released in the atmosphere upon the fuel’s combustion. The chlorination of methane: Step 1 “Initiation” UV radiation causes Cl2 to undergo homolytic ssion forming free radicals Step 2 “Propagation” Free radicals collide with methane molecules and remove a hydrogen atom: fi The methyl free radical reacts with a chlorine molecule: Step 3 “Termination” 2 free radicals collide forming a molecule; there are 3 possible products:  Check list Dr. Muhammed Gamal Topic 4 Yes Avg. No 1. What is the Definition of Hydrocarbons. 2. Understanding What does Saturated and Unsaturated refer to. 3. Understanding How many Types of Formula you should know: a. General Formula b. Molecular Formula c. Empirical Formula d. Structural Formula e. Displayed Formula (Full Structural Formula “Showing all Bonds”) f. Skeletal Formula 4. What is the Definition of Functional Group. 5. What is the Definition of Homologous Series. 6. Understanding the Properties of Homologous Series. 7. Understanding How to count the Prefix of the Compounds. 8. Understanding How to add the suffix of the Compounds. 9. Understanding How to name any Organic Compound. 10. Understanding that the Sum of all position numbers should be the Smallest. 11. What is the Definition of Structural Isomerism, and its Types: a. Chain Isomerism b. Functional group Isomerism c. Positional Isomerism 12. Understanding How to use 2n-4+1 to get the isomers of Alkanes. 13. Understanding the Difference between Cycloalkanes and Alkenes. 1  Check list Dr. Muhammed Gamal 14. Describe all Types of Reactions: a. Addition Reaction b. Substitution Reaction c. Oxidation Reaction d. Reduction Reaction e. Polymerisation Reaction 15. Understanding the Two types of Bond Breaking (Fission): a. Homolytic Fission b. Heterolytic Fission 16. Understanding The Structure of Alkyl Groups and their General Formula. 17. Understanding the Meaning of Crude Oil. 18. Understand the Processes of converting the Crude oil into Fuel: a. Fractional Distillation b. Cracking c. Reforming 19. Which is property in Organic Compounds used in using the Fractional Distillation. 20. What the Difference between the upper and lower Molecules in the Fractionating Column: a. Boiling Point b. Colour c. Passage of Light d. Viscosity 21. What are the Products and Conditions used in Cracking. 22. What is the mean of Reforming of Straight-chain Organic Compounds, and its Conditions. 23. What are the Benefits of Reforming of Straight-chain Organic Compounds. 24. What is the Difference between the Complete and Incomplete Combustion. 25. What are the Emission Problems associated with Incomplete Combustion. 2  Check list Dr. Muhammed Gamal 26. Determining the States of Matter of each Organic Compounds. 27. Understanding the Consequences of Emission of Oxides of Sulfur. 28. Understanding the Consequences of Emission of Oxides of Nitrogen. 29. Understanding How the Acid Rains are formed and its Consequences. 30. Understanding What are Catalytic Converters needed for and its Conditions. 31. What do greenhouse gases contribute to. 32. What is the meaning of the Carbon Neutrality. 33. What is the Difference between Bioethanol and Ethanol. 34. Understanding How to use Hydrogen fuel instead any other fuels. 35. What is the difference between Biofuels and Natural Gas: a. Land use b. Yield c. Manufacture/Transport d. Sustainability 36. What is the Condition needed for Free-Radical Substitution. 37. Understanding which Bond Fission used in Free-Radical Substitution. 38. Understanding the 3 Processes in Free-Radical Substitution Reaction: a. Initiation b. Propagation c. Termination 39. What product indicates that the mechanism had that sequence. 40. Understanding Why adding excess of Chlorine is not useful in Free-Radical Substitution. 41. Understanding What is the Difference between Hazards and Risks. 3 Unit 1 Topic 5 …………..Summary Menna Mohammed Badawy Key points: Sigma bonds: are single bonds involving the overlapping of 2 s-orbitals or 2 p-orbitals overlapping end-on-end, this results in a single region of overlap. Pi bonds: 2 p-orbitals that are parallel to one another overlap sideways resulting in 2 regions of overlap. (Double bond) Geometric isomers: When the atoms or groups are attached at different position on opposite sides of the C=C Bond A molecule could posses geometric isomerism due to: The Restricted rotation around the double bond Different Alkyly groups/atoms on either side of the double bond Alkenes’ general formula: CnH2n Each carbon can only form up to 4 bonds, so count carefully when drawing a structure; start by drawing all the carbon atoms, add the double bond(s) and nally add the hydrogens. Sigma-bond forms rst then a pi-bond; the following diagram represents the formation of those 2 bonds: In a pi bond the electrons are further away from the carbon atoms; unlike the sigma bond, so there is less control over the electron cloud. How to name compounds using E/Z naming system: E/trans: When the high priority atoms/groups are across one another Z/Cis: When the high priority atoms/groups are on the same side Example: fi fi Addition Reactions: Hydrogenation: addition reaction in which hydrogen is added to an alkene, under high temperature and the presence of a nickel catalyst. Halogenation: involves the addition of a halogen molecule to produce dihalogenoalkanes. Hydration: involves the addition of a water molecule to produce an alcohol. Requires heating the alkene with steam and a phosphoric acid catalyst. Oxidation to diols: Oxidising agent such as acidi ed potassium manganate reacts with the alkene forming a compound with 2 –OH groups. Mechanism of addition reactions: 1)Addition of Hydrogen hallides An electrophilic addition reaction that involves the heterolytic ssion of the H-Br molecule. fi fi Oppositely charged ions attract forming a new covalent bond; Br shares a lone pair with the carbocation. Addition reactions sometimes produce a compound in a greater portion than the other; known as the major product. In the H-Br addition reaction; the major product is the one resulting from a secondary carbocation since they are more stable that primary carbocations. If a tertiary carbocation is involved then it’s more stable. 2)Electrophilic addition of Halogens The approaching halogen molecule bonding electrons are repelled by the electron cloud of the pi-bond, causing the molecule to become polar 3)Addition polymerization: Monomer: small molecules that combine together forming a polymer Repeat unit: the set of atoms that are joined together in large number to produce the polymer structure In order to from a polymer the alkene double bond is broken and the alkene molecules are joined together: Managing polymer waste: Traditional plastic is non-biodegradable; so it doesn’t breakdown. Incineration: Burning polymer waste; this gets rid of the polymer but produces toxic gasses which pollutes the atmosphere and release carbon dioxide, a greenhouse gas.  Check list Dr. Muhammed Gamal Topic 5 Yes Avg. No 1. Memorising the meaning of Unsaturated Hydrocarbons. 2. Knowing the General Formula of Alkenes. 3. Understanding the General Formula of Cycloalkenes. 4. Memorising the Types of Bonds in a Double Bond. 5. Understanding the Difference between Sigma Bond and Pi Bond. 6. Understanding How to Draw Sigma and Pi Bonds of a Double Bond. 7. Knowing the Difference between Structural isomerism and Stereoisomerism. 8. Understanding the Conditions that must settle in a Compound to have a Geometric isomerism. 9. Understanding Why Alkanes cannot have a Geometric isomerism. 10. Understanding that each Double Bond has 2 Geometric isomers if the Conditions are Found. 11. Understanding When to identify an Alkene having Cis/Z Compound or Trans/E Compound depending on Highest Atomic Number on each Side of the Double Bond. 12. Understanding the Types and the Definition of Addition Reactions, and its Conditions: a. Hydrogenation b. Halogenation c. Hydration d. Hydrogen Halides e. Oxidation to Diol 13. Memorising the Test for Alkenes and Cycloalkanes. 1  Check list Dr. Muhammed Gamal 14. What is the Definition of Electrophilic Addition. 15. What is the Definition of Electrophile. 16. Understanding the Mechanism of Electrophilic Addition: a. Hydrogen Halides b. Halogens 17. Understanding the meaning of Major and Minor Products from Electrophilic Addition that refers to Primary, Secondary and Tertiary Carbocations. 18. Understanding How to Convert between Addition Polymerisation and Monomers. 19. Understanding the meaning of Non-Biodegradable polymers and Biodegradable Polymers. 20. Understanding the Problems associated with production of Non-Biodegradable polymers. 21. Understanding the Solutions to minimise the Polymer wastes. 22. What is the Definition of Incineration and its Consequences. 23. What are the Advantages and Disadvantages of using Biodegradable Polymers. 2  Essential Equations Dr. Muhammed Gamal Unit 1 Essential Equations Reaction Name Equation Conditions Reaction Identification First Ionisation Energy All atoms and Ions Converting from an Atom Ag(g) → Ag+(g) + e- should be Gaseous into a Positive Ion Second Ionisation Energy Converting from a Singly All Ions should be Ag+(g) → Ag+2(g) + e- Charged Positive ion into a Gaseous Doubly Charged Positive Ion Mass Spectrometry Ni(g) + e- → Ni+(g) + 2e- All atoms and Ions Bombarding a Species with Cl2(g) + e- → Cl2+(g) + 2e- should be Gaseous an electron to be Charged Homolytic Fission ↷ Producing 2 Free Radicals Cl – Cl → 2Cl ↷ - with unpaired electron Heterolytic Fission ↷ Producing 2 oppositely H – Cl → H+ + Cl- - charged Ions Complete Combustion Converting Organic CH4 + 2O2 → CO2 + 2H2O - Compounds into Carbon dioxide and Water Vapour Incomplete Combustion Converting Organic CH4 + 3/2 O2 → CO + 2H2O - Compounds into Carbon Monoxide and Water Vapour 1  Essential Equations Dr. Muhammed Gamal Reaction Name Equation Conditions Reaction Identification Free Radical Substitution Converting Alkane into CH4 + Cl2 → CH3Cl + HCl Ultraviolet Light Halogenoalkane Initiation Producing 2 Free Chlorine ↷ Cl ↷ Cl → 2Cl Ultraviolet Light Radicals with unpaired electron Propagation Cl + CH4 → HCl + CH3 Bombarding of Free - CH3 + Cl2 → CH3Cl + Cl Radicals with Species Termination Cl + Cl → Cl2 Two Free Radicals Joined Cl + CH3 → CH3Cl - Together CH3 + CH3 → C2H6 Further Substitution Converting Chloroalkane to CH2Cl + Cl → CH2Cl2 - dichloroalkane or further Cracking High Temperature Converting Long Chain C10H22 → C5H12 + C5H10 + Alkane into Alkene(s) Zeolite Catalyst Reforming Heat Converting Straight Chain + Compounds into Branched Platinum Catalyst or Cyclic Compounds Hydrogenation Heat under Pressure Converting Alkene into H2C=CH2 + H2 → CH3CH3 + Alkane Nickel Catalyst 2  Essential Equations Dr. Muhammed Gamal Reaction Name Equation Conditions Reaction Identification Bromination Converting Alkene into CH3CH=CH2 + Br2 → CH3CHBrCH2Br - Dibromoalkane Mechanism of H H H H — Electrophilic Addition C—C H C—C + Br - — (Hydrogen Bromide) — H H H δ+ H H — δ- Carbocation Br - Converting Ethene into H H H H — — — Bromoethane H — C—C H — C—C — H — — — H H Br - H Br Mechanism of H H H H — Electrophilic Addition C—C H C—C + Br - — — (Bromination) H H H δ+ Br Br — δ- Carbocation Br - Converting Ethene into H H H H — — — 1,2-dibromoethane H C—C H C—C H — — — — — — H Br Br - Br Br 3  Essential Equations Dr. Muhammed Gamal Reaction Name Equation Conditions Reaction Identification Mechanism H of H H H — High C—C H C—C + OH- — Electrophilic Addition — H H H Temperature (Hydration) δ- δ+ H O H (300ºC) Carbocation δ+ H + 60 atm Converting Ethene + H H H H — — — into Phosphoric H C—C H C—C H — — — Ethanol Acid — — — H H OH- H OH Test for Alkene (Adding Bromine Water) CH3CH=CH2 + Br2(aq) → CH3CH(OH)CH2Br - Converting Ethene to Bromoethenol Aqueous Oxidation to Diol Potassium Manganate CH3CH=CH2 + [O] + H2O → CH3CH(OH)CH2OH (VII) Converting Alkene into Diol + Dilute H2SO4 Addition Polymerisation High Repeat Unit H H Temperature H H n C C C C + Converting Alkene to H X H X High Pressure n Poly(Alkene) Monomer Polymer + Catalyst 4