NOTES Index Part 1 The periodic table - Atomic Structure. States of matter. (10 -22) Chemical formula of compound. (23-28) Bonding. (29 -34) Material structure. (35-42) Elements. (43-61) Colours . (62 -63) Types of Chemical Reaction (64 -79) Energy and Chemistry. (74-80) Acids and Alkalis. (81-90) Salt Preparation. (91 - 107) Mixtures Separation and purification. practical work (1) (102 -108) (109 - 117 ) Atomic Structure Matter is anything that has mass and takes up space (volume). Chemistry is the study of how matter behaves and how one kind of substance can be changed into another. Matter consists of molecules. A molecule is a group of atoms held together by covalent bonds . A molecule in turn consists of atoms. An atom is the smallest particle of an element that takes part in a chemical reaction without being changed. Matter can be classified into pure substances and mixtures. More than one substance mixed in any ratio. impurities Examples: air, sea water and petroleum. An element is a substance that consists of one type of atoms that can not be splitted into simpler substances. bonded together. Examples: water, Carbon dioxide, Sodium chloride. The periodic table - Atomic Structure. Amani Hamdi Page 1 Each element has a symbol, for example: Aluminum Phosphorous Carbon Calcium Al P C Ca Sulfur S Potassium K Chlorine Cl Iron Fe Sodium Silicon Copper Nitrogen Na Si Cu N Elements can be classified into metals and non-metals. Elements are monoatomic, diatomic and polyatomic. Elements Monoatomic The molecule of the element is formed of one atom. Examples: Na, K, Al, Ca, and all noble gases. Diatomic The molecule consists of two atoms. Examples: H2, N2, O2 and halogens (F2, Cl2, Br2, I2 and At2) Polyatomic The molecule consists of more than two atoms. Examples: P4 , S8 A Compound is a substance that is formed of different elements that combine together in a fixed ratio by chemical bonds. Although there is a limited number of elements, there are millions of compounds. The periodic table - Atomic Structure. Amani Hamdi Page 2 Atomic Structure An atom is the smallest particle of a substance that can share in a chemical reaction without being changed. It consists of a nucleus, and a cloud of particles called electrons that rotate around the nucleus. 6 electrons outside nucleus nucleus contains 6 protons and 6 neutrons The nucleus is a cluster of two sorts of sub- particles, protons and neutrons. All the particles in an atom are very light. Their mass is measured in atomic mass units (a.m.u), rather than grams. Protons and electrons also have an electric charge. Particle in atom Proton Neutron Electron Relative mass 1 a.m.u 1 a.m.u 1/1843 ~ 1/2000 a.m.u Relative charge Positive change (+1) zero Negative change (-1) The protons are positively charged, the electrons are negative and the neutrons are neutral (have no charge). The electrons rotate around the nucleus in energy levels (electron shells). Each shell can hold a limited number of electrons. The 1st shell The 2nd shell The 3rd shell 2 electrons 8 electrons 18 electrons Atomic number (Proton number): It is the number of protons in the nucleus of an atom. Mass number ( nucleon number) : It is the sum of protons and neutrons in the nucleus of an atom. The electrons in an atom have almost no mass. So the mass of an atom is nearly due to its protons and neutrons The periodic table - Atomic Structure. Page 3 Amani Hamdi Shorthand form of an atom: An atom can be described in a short way, using its symbol, its proton number and its mass number in a certain order as the following: Z A The proton number Symbol of element The mass number Note that: The atomic no. = the no. of protons = the no. of electrons in an atom. The mass number = no. of protons + no. of neutrons. So the no. of neutrons = the mass no. – Atomic no. Example: 11 23 Na This is a Sodium atom. That has 11 protons, 11 electrons and 12 neutrons (23-11=12) Valency: Nobel gases are stable or uncreative because they have completely filled outermost shell. They have eight electrons in their outermost shell except Helium that has 2 electrons (as it has only 1 shell that is completed by only 2 electrons). All atoms tend to have the same electronic structure as noble gases. This can be done by losing, gaining or sharing electrons from the outermost shell of the atom during a chemical reaction. The number of electrons lost, gained or shared by an atom is called the valency of an element. Metals tend to lose the outermost shell electrons. Non-metals tend to gain or share electrons. Transition elements have more than one valency. The periodic table - Atomic Structure. Page 4 Amani Hamdi The following table shows the Valencies of some elements: Element Sodium Chlorine Magnesium Oxygen Aluminum Nitrogen Carbon Silicon Symbol Electronic configuration 11Na 2, 8, 1 2, 8, 7 2, 8, 2 2, 6 Change of formed ion Mono valent 1+ Mono valent 1Divalent 2+ Divalent 2- 2, 8 , 3 2,5 2, 4 2, 8, 4 Trivalent Trivalent Tetravalent Tetravalent 17Cl 12Mg 8O 13Al 7N 6C 14Si Valency 3+ 3- Notes loses 1é gains 1 é loses 2 é gains 2é loses 3é gains 3é Tend to share 4é Atoms of metals have 1, 2 or 3 electrons in the outermost shell except for Hydrogen, Helium and Boron, they are non-metals. Atoms of non-metal have 4, 5, 6 or 7 é in their outermost shell. Nobel gases have completely filled outermost shell (atoms of noble gases have 8 electrons in the outermost shell except for Helium it has only 2 electrons), so they have a valency of zero. Metals tend to lose electrons. Non-metals that have 4 electrons tend to share electrons. Non-metals that have 5,6 or 7 electrons tend to gain or share electrons. The periodic table - Atomic Structure. Page 5 Group 0 Group 7 Group 6 Group 5 Transition Metals He 3+ 4± 3- 2- 1- 0 Group 4 Group 2 Period 3 1+ 2+ Group 1 Period 2 H Valency Group 3 Period 1 Amani Hamdi Isotopes Isotopes: atoms of the same element which have the same proton number but different nucleon number. Examples: H- 1 , H -2 , H -3 Cl- 35 , Cl-37 U -235 , U -238 The atoms of the isotopes have the same number of protons but different numbers of neutrons in their nuclei. Isotopes of an element have the same chemical properties because they have the same electron structure and so same olinumber of electrons in the outermost shell. The relative atomic mass (Ar) is the average mass of naturally occurring atoms of the same element on a scale where carbon-12 atom has a mass of exactly 12 units. The relative atomic mass, Ar, of an element is calculated from: the mass numbers of its isotopes the abundance of these isotopes The periodic table - Atomic Structure. Amani Hamdi Page 6 For e.g : Chlorine Chlorine naturally exists as two isotopes, Cl (chlorine-35) and Cl (chlorine-37). The abundance of chlorine-35 is 75% and the abundance of chlorine-37 is 25%. In other words, in every 100 chlorine atoms, 75 atoms have a mass number of 35, and 25 atoms have a mass number of 37. To calculate the relative atomic mass, Ar, of chlorine: Ar=total mass of atoms / total number of atoms = (35 x 75) + (37 x 25) (75 + 25 ) Ar = 2625 + 925 100 = 3550 100 Ar=35.5 (to 1 decimal place) Notice that the answer is closer to 35 than it is to 37. This is because the chlorine-35 isotope is much more abundant than the chlorine-37 isotope. Amani Hamdi The periodic table - Atomic Structure. Page 7 The Periodic Table Scientists put all the elements into groups according to their atomic number in an ascending order as follows. Hydrogen comes first, as its proton number (atomic no.) is 1. The table has eight vertical groups of elements, plus a block of transition metals. The last group is the group zero. Their atoms have completely filled outermost shells. Some groups have special names: Group 1 is called Alkali Metals. Group 2 is called Alkaline Earth Metals. Group 7 is called Halogens. Group 0 is called Noble gases. Transition metals lie in the middle of the periodic table between group 2 and 3 beginning from the 4th period. They include most hard and dense metals. The horizontal rows in the periodic table are called periods. There are several trends as you move across the period, for example; the elements go from metals to non-metals. Some substances between metals and non-metals like Silicon are like metals in some ways and non-metals in other ways. They are called Metalloids. Hydrogen stands on its own in the periodic table. This is because it has one outermost electron, like the group 1 metals, but unlike them it is a gas and it usually reacts like non-metals. Not all elements occur naturally, some are artificial (synthesized by scientists) during nuclear reactions. Most of these elements lie at the bottom block of the periodic table. All isotopes of artificial elements are radioactive. This means their atoms have unstable nuclei which break down giving out radiation. The periodic table - Atomic Structure. Page 8 Amani Hamdi The Periodic Table of Elements Group I III II Key 3 4 atomic number IV V VI VII VIII 1 2 H He hydrogen helium 1 4 5 6 7 8 9 10 Li Be atomic symbol B C N O F Ne lithium beryllium Name boron carbon nitrogen oxygen fluorine neon 7 9 relative atomic mass 11 12 14 16 19 20 11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar sodium magnesium aluminium silicon phosphorus sulfur chlorine argon 23 24 27 28 31 32 35.5 40 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr potassium calcium scandium titanium vanadium chromium manganese iron cobalt nickel copper Zinc gallium germanium arsenic selenium bromine krypton 39 40 45 48 51 52 55 56 59 59 64 65 70 73 75 79 80 84 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe rubidium strontium yttrium zirconium niobium molybdenum technetium ruthenium rhodium palladium silver cadmium indium tin antimony tellurium iodine xenon 85 88 89 91 93 96 – 101 103 106 108 112 115 119 122 128 127 131 55 56 57–71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba lanthanoids Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn caesium barium hafnium tantalum tungsten rhenium osmium iridium platinum gold mercury thallium lead bismuth polonium astatine radon 133 137 178 181 184 186 190 192 195 197 201 204 207 209 – – – 87 88 89–103 104 105 106 107 108 109 110 111 112 actinoids 114 116 Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Fl Lv francium radium rutherfordium dubnium seaborgium bohrium hassium meitnerium darmstadtium roentgenium copernicium flerovium livermorium – – – – – – – – – – – – – lanthanoids actinoids 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu lanthanum cerium praseodymium neodymium promethium samarium europium gadolinium terbium dysprosium holmium erbium thulium ytterbium lutetium 139 140 141 144 – 150 152 157 159 163 165 167 169 173 175 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr actinium thorium protactinium uranium neptunium plutonium americium curium berkelium californium einsteinium fermium mendelevium nobelium lawrencium – 232 231 238 – – – – – – – – – – – The volume of one mole of any gas is 24 dm 3 at room temperature and pressure (r.t.p.). Page 9 States of matter There are three physical states of matter Solid, liquid and gas. The individual particles of the substance are the same in each state. It is their arrangement that is different. Physical State Volume Density Shape Fluidity Solid Fixed volume High Definite shape Doesn’t flow Moderate to high No definite shape ,takes the shape of the container Flows easily Low No definite shape, takes the shape of the container Flows easily Liquid Fixed volume Gas No Fixed volume expands to fill the container Note: liquids and gases are called fluids. States of Matter Amani Hamdi Page 10 State Separation of particles Solid Closely packed Liquid Gas Arrangement of particles Motion of particles Regular Vibration about fixed point. Fairly closely packed Random Move about and slide past each other. Far apart, Random Move about very fast in random motion. The volume of a gas at a fixed temperature can be easily reduced by increasing the pressure on the gas. They are easily compressed. Liquids are only slightly compressible, and the volume of a solid is unchanged by changing the pressure. Particles of gases collide with each other and with the sides of the container (This is how pressure is created inside a sealed container ) States of Matter Amani Hamdi Page 11 Melting Solid Heat Energy" Liquid "by gainig Heat Energy" Freezing by losing heat by losing heat Solid Sublimation Solid Melting: is the change of a substance from solid state to liquid state. Melting point: is the temperature at which a solid begins to melt. The melting point is the same value as the freezing point. For example, the melting point and freezing point of pure water takes place at 0oC. States of Matter Amani Hamdi Page 12 Sublimation: is the direct change from solid to gas without passing through liquid state. Few solids can sublime such as, dry ice (solid carbon dioxide) and iodine sublimes producing a purple vapour, then solidifies again on cold surface to return back a dark grey solid. If a liquid is left with its surface exposed to the air, it evaporates. Evaporation: is the process in which a liquid changes into gas. Evaporation takes place from the surface of the liquid. The larger the surface area, the faster the liquid evaporates. The warmer the liquid is, the faster it evaporates. So evaporation is affected by surface area and temperature. At a certain temperature the liquid becomes hot enough for a gas to form within the liquid and not just on the surface. Bubbles of gas appear inside the liquid. This process is known as boiling. Boiling is a condition under which gas bubbles are able to form within a liquid and gas molecules escape from the body of a liquid, not just from its surface. Boiling point: the temperature at which liquid boils, when the pressure of gas created above the liquid equals atmospheric pressure. A volatile liquid is one which evaporates easily and has a relatively low boiling point. Ethanol (boiling point 78°C) is more volatile than water (boiling point 100°C). States of Matter Amani Hamdi Page 13 Evaporation Boiling it occurs over a range of temperatures it occurs only at the surface of liquids does not produce bubbles occurs at specific temperature known as boiling point takes place throughout the bulk of the liquid. produces lots of bubbles The reverse of evaporation is Condensation. Condensation: is the change of a vapour or a gas into a liquid. Condensation is usually brought about by cooling. It is possible at normal temperature to condense a gas into a liquid by increasing pressure, without cooling. Pure substances: A pure substance consists of only one substance. There is no contaminating impurities. A pure substance melts and boils at definite temperatures. The melting point and boiling point can be used to test for purity of a substance, also used to check the identity of an unknown. Amani Hamdi States of Matter Page 14 A substance’s melting point and boiling point in relation to room temperature (taken as 20°C) determines whether it is usually seen as a solid, liquid or a gas. (State of matter) If the melting point is below 20°C and the boiling point is above 20°C, the substance will be a liquid at room temperature. If the melting point is above room temperature the substance is a solid. If the boiling point is below room temperature, the substance will be a gas at room temperature. The melting points of some common chemical substances: Physical state at room Substance temperature(20◦C) Oxygen Gas Nitrogen Gas Ethanol(alcohol) liquid Water liquid Sulfur Solid Common salt(Sodium Solid Chloride) Copper Solid Carbon dioxide Gas Melting point/◦C Boiling point/◦C -219 -210 -117 0 115 -183 -196 78 100 444 801 1465 1083 -78 2600 (sublimes) States of Matter Amani Hamdi Page 15 The figure at the right shows how the temperature changes when a sample of the solid Naphthalene (a single pure substance) is heated steadily. -The solid melts sharp at 80°C. While the solid is melting the temperature stops rising. It will begin to rise again when all the Naphthalene has melted. -If you continue to heat the liquid until its boiling point, the temperature stays the same until all the liquid has boiled. heating curve for wax, which is a mixture of substances, -The shows the solid wax melting over a range of temperatures. Amani Hamdi States of Matter Page 16 Kinetic theory Thermal energy (heat energy) which is converted into kinetic energy . This is basic of kinetic theory of matter Heating a solid causes its particles to vibrate more and as the temperature increases , they vibrate so much that the solid expands , until the structure breaks and the solid melts On further heating , the liquid substance expands more and some particles at the surface gain sufficient energy to overcome the intermolecular forces and evaporate . when the boiling point temperature is reached , all the particles gain enough energy to escape and the liquid boils. These changes in state can be shown on a graph called a heating curve . Amani Hamdi States of Matter Page 17 cooling down a gas has the reverse effect and this would be called a cooling curve when the gas is heated , the particles lose thermal energy so the kinetic energy decreases and the particles slow down. States of Matter Amani Hamdi Page 18 Diffusion in fluids The idea that fluids (liquids and gases) are made up of moving particles helps us to explain processes involving diffusion. -Dissolving Potassium Manganate (VII) crystals is placed at the bottom of a dish of water. It is then left to stand. At first the water around the crystals becomes purple as the solid diffuses. Particles move off the surface of the crystal into the water. The crystal dissolves completely and the whole solution becomes purple. The particles from the solid become evenly spread through the water. -The spreading of the solute particles throughout the liquid is an example of diffusion. -Diffusion in solution is also important when the solute is a gas. This is important in breathing. Diffusion contributes to the movement of oxygen from the lungs to the blood, and of carbon dioxide from the blood to the lungs. States of Matter Amani Hamdi Page 19 Diffusion of gases -A few drops of liquid bromine are put into a gas jar and covered with another gas jar as shown. After short time the jar becomes full of brown gas. Bromine gains energy from the surrounding The intermolcular forces between particles of liquid bromine get weaker . Bromine Vaporizes . Particles of bromine and air collide with each other . Bromine diffuses from the area of high concentration to that of low concentration . Definition of Diffusion: Is the process by which different fluids mix as a result of random motion of their particles. Diffusion involves the movement of particles from a region of higher concentration towards a region of lower concentration. Diffusion does not take place in solids. Diffusion in liquids is much slower than in gases. Not all the gases diffuse at the same rate. States of Matter Amani Hamdi Page 20 -The Ammonia and the Hydrogen chloride acid fumes react when they meet producing a white smoke ring of Ammonium Chloride. HCl(g) + NH3(g) NH4Cl(s) -The fact that the ring is not formed halfway along the tube shows that ammonia is a lighter molecule than hydrogen chloride acid fumes as the ring is formed nearer to HCl rather than NH3. -To show which gas is lighter, we have to calculate the relative molecular mass (RMM or Mr) of each molecule which is the sum of the atomic mass of each atom. RMM of HCl = 1 + 35.5 = 36.5 RMM of NH3 = 14 + (3 x 1) = 17 NH3 has less RMM than HCl, so Ammonia diffuses at faster rate. States of Matter Amani Hamdi Page 21 -The speed of gas atoms or molecules are high. These particles collide very frequently with other particles in the air (many millions of collisions per second). So their path is not direct. See the figure at the right. -These very frequent collisions slow down the overall rate of diffusion from one place to another. -The amount of individual gas molecules or atoms in the air can not be seen as the particles are too small. N .B . Heavier particles move more slowly than lighter particles at the same temperature. Larger molecules diffuse more slowly than smaller molecules. The pressure of a gas is the result of collisions of the fast-moving particles with the walls of the container. The average speed of the particles increases with an increase in temperature. States of Matter Amani Hamdi Page 22 Chemical Formula Formula of ionic compounds: Ionic compounds are formed between metal atoms and non-metal atoms. Metal atoms form positive ions (cations) by losing electrons. Non – metal atoms form negative ions (anions) by gaining electrons. The positive and negative ions attract each other forming an ionic bond. The overall structure must be neutral i.e. positive and negative charges must balance each other. Examples: Chemical formula of Magnesium Chloride The ions present: Mg2+ ClClTotal charge: 2+ 2- This is why the formula is written as MgCl2 The size of the charge on an ion is a measure of its valency. Amani Hamdi Chemical Formula Page 23 Ionic Compounds Binary Polyatomic Formula of binary ionic componds : Examples : Formula of Aluminum Oxide Write down the correct symbol. Al Write down the charges of the ions. Reverse the Valencies and skip the charges. O Al2O3 Formula of Calcium Oxide Write down the correct symbol. Ca Write down the charges of the ions. Reverse the Valencies and skip the charges. O Ca2O2 Formula of Manganese (IV) Oxide CaO Simplify Write down the correct symbol. Mn Write down the charges of the ions. Reverse the Valencies and skip the charges. O Mn2O4 MnO2 Simplify Amani Hamdi Chemical Formula Page 24 Formula of Iron (III) Oxide Write down the correct symbol. Fe Write down the charges of the ions. Reverse the Valencies and skip the charges. O Fe2O3 Formula of Iron (II) Oxide Write down the correct symbol. Fe Write down the charges of the ions. 2+ Reverse the Valencies and skip the charges. O 2Fe2O2 Simplify N.B.Transition elements have more than one valency. Formula of ionic compounds containing polyatomic ions: Follow the same rules to write the formula. You have to study and memorize the formula and the charges of the polyatomic ions. Monovalent Radicals Divalent Radicals Hydroxide(OH-) Trivalent Radicals Carbonate(CO3-2) Nitrate(NO3-) Nitrite(NO2-) Sulfate(SO4-2) - Hydrogen Carbonate(HCO3 ) Ammonium(NH4+) Phosphate(PO4-3) Sulfite(SO3-2) Amani Hamdi Chemical Formula Page 25 Examples: Formula of Sodium Carbonate Write down the correct symbol. Na Write down the charges of the ions. 1+ Reverse the Valencies and skip the charges. CO3 Na2CO3 Formula of Ammonium Phosphate Write down the correct symbol. NH4 PO4 Write down the charges of the ions. Reverse the Valencies and skip the charges. (NH4)3PO4 Naming Ionic Compounds The positive part (metal ion) is the prefix. The negative part (non-metal ion or the negative polyatomic ion) is the suffix. For non-metal ions, the suffix is ide e.g.Calcium Chloride, CaCl2 For negative polyatomic ions, the suffix is either ite or ate E.g. Sodium Nitrite, NaNO2 Sodium Nitrate, NaNO3 Hint: ite has less oxygen than ate. Roman figures are used to indicate the Valencies of transition metals, as they have more than one valency. E.g. Iron(II)Oxide, FeO Iron(III)Oxide, Fe2O3 Amani Hamdi Chemical Formula Page 26 Naming Covalent Compounds Covalent compounds are formed between non-metal atoms. The prefix is the name of the first element. The suffix is the name of the second element, ends in ide. Prefixes like (Mono,Di,Tri…….) are used to indicate the number of atoms of each element. Mono is not used for the first element e.g. NO2 Nitrogen dioxide, i.e. we do not say mono Nitrogen dioxide. When the addition of the Greek prefix places two vowels adjacent to one another, the “a” (or the “o”) at the end of the Greek prefix is usually dropped; e.g., “nonaoxide” would be written as “nonoxide”, and “monooxide”would be written as “monoxide”. The “I” at the end of the prefixes “di-“ and “tri” are never dropped. Prefix MonoDiTriTetra Penta- Number Indicated 1 2 3 4 5 Prefix HexaHeptaOctaNonaDeca- Number Indicated 6 7 8 9 10 Amani Hamdi Chemical Formula Page 27 Writing covalent compounds from their names: Carbon tetrachloride: CCl4 Dinitrogen tetroxide: N2O4 Writing covalent names from their formulas: I4O10 Tetraiodinedecoxide BrCl3 Bromine trichloride Complete the following table: Formula N2F6 CO2 Name Dinitrogen hexafluoride Silicon tetrafluoride CBr4 Nitrogen trichloride P 2S 3 Carbon monoxide NO2 Carbon tetrachloride PCl3 Amani Hamdi Chemical Formula Page 28 Bonding Atoms combine with other atoms to be able to achieve a stable arrangement of electrons like inert gases that have completely filled outermost shell ( Octet rule). Atoms combine by Ionic bonds or Covalent bonds Ionic bonds Ionic bond is formed between metal and non-metal atoms. A metal atom loses one or more electrons from its outer most shell forming a Positive ion . A non-metal atom gains one or more electrons forming a negative ion Electrostatic attraction takes place between the +ve and –ve ions forming an ionic compound with an ionic bond. Ionic bonding is the strong electrostatic attraction force between oppositely charged ions. Examples: 1. Formation of Sodium Chloride (NaCl) Bonding Amani Hamdi Page 29 2. Formation of Magnesium Fluoride (MgF2) Covalent bonds Covalent bonds take place between non-metal atoms by sharing electrons. Covalent bonding is a Chemical bonding formed by the sharing one or more pairs of electrons between two atoms. Covalent bonds Single covalent bond A covalent bond that is formed by sharing two electrons, one from each nonmetal atom. Double covalent bond A covalent bond that is formed by sharing four electrons, two from each non-metal atom. Bonding Triple covalent bond A covalent bond that is formed by sharing six electrons, three from each non-metal atom. Amani Hamdi Page 30 Examples for single covalent compounds: C + Atomic no. 6 1. A Methane molecule CH4 4H 1 - Electronic 2,4 configuration H H H x C xH H H 2. A water molecule H2O oX O xH x x H H Atomic no. 1 - Electronic 1 configuration xx H x H oX Hx xx O H 8 2,6 1 1 O x 3. A Hydrogen molecule H2 H Atomic no. - Electronic configuration H xH x x x H: H 1 1 1 H Bonding H 1 1 H H Amani Hamdi Page 31 4. A Hydrogen chloride molecule HCl H Atomic no. - Electronic configuration 5. A chlorine molecule Cl2 Cl Atomic no. - Electronic configuration 6. A methanol molecule CH3OH Atomic no. - Electronic configuration 1 1 17 2,8,7 C 4H Bonding 17 2,8,7 Cl 17 2,8,7 17 2,8,7 Cl 1 1 O 8 2,6 Amani Hamdi Page 32 7. An ethanol molecule C2 H5 OH 2C Atomic no. - Electronic configuration 17 2,8,7 6H 1 1 O 8 2,6 Example of double covalent bond in molecules: 8. An Oxygen molecule O2 O Atomic no. 8 - Electronic 2,6 Configuration O O 8 2,6 O O 9. An ethene molecule C2 H4 O 2C Atomic no. - Electronic configuration 17 2,8,7 Bonding 4H 1 1 Amani Hamdi Page 33 Example of triple covalent bond in molecules: N A nitrogen molecule N2 Atomic no. 7 Electronic 2,5 configuration N 7 2,5 N Comparison between ionic and covalent compounds: Properties of typical ionic Compounds Properties of simple covalent Compounds They are crystalline solids at room temperature. They are often liquid or gases at room temperature. They have high melting and boiling points. They have low melting and boiling points. Volatile Non-Volatile They are often soluble in water (not usually soluble in organic solvents, e.g. ethanol, methylbenzene). They are soluble in organic solvents such as ethanol or methylbenzene (very few are soluble in water). They conduct electricity when molten or dissolved in water (not when solid). They do not conduct electricity. Bonding Amani Hamdi Page 34 Material structure There are four types of solid physical structure Metallic Structure (Metal crystals) Metal atoms have relatively few electrons in their outer most shells. When they are packed together each metal atom loses its outer electrons into “sea” of free electrons (or mobile electrons). The atoms are no longer electrically neutral; they became ions because they have lost electrons. Therefore the structure of a metal is described as positive ions packed in a lattice surrounded by a sea of mobile electrons (delocalized electrons). Material Structure Amani Hamdi Page 35 Delocalized electrons: Electrons those are not restricted to orbiting one positive ion. Lattice: Three dimensional regular arrangements of atoms, molecules or ions in a crystalline solid. Metallic Bonding: The electrostatic attraction force between the sea of mobile electrons and the regular array of positive metal ions within a solid metal. Properties of metals due to their structure: 1. High density: metals are high in density due to their close packing of ions. In some metals the ions are less closely packed. These metals have the lowest densities of all metals for example the alkali metals. 2. Metals are good conductors of heat due to the movement of the mobile free electrons. 3. Metals are good electric conductor due to the presence of free electrons. 4. Metals are malleable because the layers of identical ions can slide over one another without breaking the structure, due to strong metallic bonding. 5. Metals are ductile (can be stretched into wires) due to the strength of the metallic bonds that is strong but not rigid (metals have a high tensile strength). Material Structure Amani Hamdi Page 36 Giant Ionic Lattice (Ionic crystals) Ionic compounds structure: lattices of positive and negative ions. The nearest neighbors of the ion are always of the opposite charge. Example: in Sodium Chloride Each Sodium ion Na+ is surrounded by six Chloride ions Cl- and each Chloride ion is surrounded by six Sodium ions Na+. The actual arrangement of the ions in the other compounds depends on the number of ions involved and on their sizes. Properties of the Ionic crystals: Ionic crystals are hard but much brittle than metallic crystals. Ionic crystals are hard because pushing one layer against another in an ionic crystal bring ions of the same charge next to each other. The repulsion forces the layers apart. Ionic lattices dissolve in water. When an ionic crystal dissolves, each ion becomes surrounded by water molecules. This breaks up the lattice and keeps the ions apart. For those ionic compounds that do not dissolve in water, the forces between the ions must be very strong. Ionic compounds can conduct electricity when dissolved in water or when they are melted because, the ions are able to move and carry the current. Material Structure Amani Hamdi Page 37 Simple Molecular Structure (Molecular Crystals) Structure : Lattices of molecules . Some non-metals (e.g. Iodine and Sulfur) and some covalent bonded compounds exist as solids with low melting points. In these crystals molecules of these elements or compounds are held together by weak intermolecular forces to form crystals that are easily broken down by heat. Properties of Simple molecular structure: Molecular structure does not conduct electricity because they have no charges. Molecular structures have low melting and boiling points. This is because there are only weak forces between the molecules (Van der Waals forces). Material Structure Amani Hamdi Page 38 Giant Molecular Crystals (Macro Molecules) This type of structure occurs either between atoms of an element such as Carbon atoms in Graphite and Diamond or in compounds such as Silica SiO2 (sand) and Germanium Oxide GeO2. Graphite structure: Van dear weal’s weal’s forces How the layers fit together One layer of graphite It is a giant molecular structure composed of carbon atoms. Carbon atoms are arranged in flat layers. Each carbon atom is bonded to other three carbon atoms by strong covalent bonds. Between these layers, there are weak forces of attraction (Van der Waals forces). Graphite is soft and feels slippery because it is layered structure with weak forces of attraction between the layers, so layers are able to slide over each other. Graphite conducts electricity due to free electrons. Note: In graphite each carbon atom uses only three electrons of the four valency electrons that are used by the layered atoms in covalent bonding. There is still an electron from each atom to move between layers, this is why graphite conducts electricity. Material Structure Amani Hamdi Page 39 Diamond Structure: C C C C C Diamond is a form of carbon. Each carbon atom is attached to other four carbon atoms by covalent bonds in a tetrahedral structure. Properties of diamond: Diamond has high melting point. Hard because the strong covalent bonds extends throughout the whole structure. Brittle due to rigidity of bonds. Diamond does not conduct electricity because there are no free electrons. Note: diamond has no free electrons as all the four valency electrons are used in bonding. Material Structure Amani Hamdi Page 40 A comparison of the properties and uses of diamond and graphite: Diamond Appearance Hardness Density Graphite Properties Uses Properties Colorless, transparent crystals that sparkle in light In jewelery and ornamental objects Dark grey, shiny solid The hardest natural substance In drill bits, diamond saws and glass-cutters Soft-the layers can slide over each other-and solid has a slippery feel More dense than graphite(3.51g/ Electrical Conductivity Uses In pencils, and as a lubricant Less dense than diamond(2.25g/ Does not conduct electricity Conducts electricity As electrodes and for the brushes in electric motors Silicon (IV) Oxide (silica) O Silicon (IV) Oxide has a giant molecular structure. Sand and quartz are examples of silica. (Silicon (IV) oxide or silicon dioxide SiO2) Si O O O Si O O O Structure of silica: Each Silicon atom is attached to four oxygen atoms by covalent bonds in a tetrahedral structure and each oxygen atom is attached to two silicon atoms. Material Structure Amani Hamdi Page 41 Properties of silicon dioxide: Hard. High melting and boiling point. Does not conduct electricity. Note: Diamond and Silicon dioxide have similar tetrahedral structure and so similar properties. Material Structure Amani Hamdi Page 42 Elements An element is a substance which cannot be further divided into simpler substances by chemical methods. It is formed of only one type of atoms. • There are 94 naturally occurring elements. Most of elements (70) can be classified as metals. • Metals are separated from non-metals in the periodic table. The non-metals are grouped into the top right region of the periodic table above the thick stepped line. • The chemical properties of metals and non-metals are different. Comparison between metals and non-metals: Metals Non-Metals 1. They are usually solids, (except for mercury, which is liquid) at room temp 1.They are solids or gases (except for bromine which is liquid) at room temp. 2. Have high melting point and high boiling point. 2.Have low melting and a boiling point. 3. All metals Conduct electricity. 3.Don’t Conduct electricity (except graphite, a form of carbon) They tend to be insulators. 4. They are good conductor of heat 4.They are poor thermal conductor. 5. They are malleable (their shape can be changed by hammering. They are ductile (can be pulled out into wires) 5.Most non-metals are brittle when solid. 6. They are grey in colour(except for gold and copper) 6.They vary in colour they have a dull surface when solid 7. They make a ringing sound when struck (They are sonorous) They are not sonorous Elements Amani Hamdi Page 43 Family profile of some groups Group I The alkali metals -Metals of group I are called alkali metals. They are the most reactive metals. They are known as alkali metals because they react vigorously with water producing Hydrogen and alkaline solution. -Alkali metals are Lithium, Sodium, Potassium, Rubidium, Cesium and the radioactive element Francium. Physical Properties of alkali metals • They are soft solids. • They have relatively low melting points. • They have low densities. Chemical Properties of alkali metals • They are highly reactive, so they are stored in oil to prevent them reacting with oxygen and water vapour in the air. • They have a valency of +1 Note: • When freshly cut with knife all these metals have a light grey, silvery surface which quickly tarnishes (becomes dull) Alkali metal trends • Reactivity increases as we go down the group. All group I reacts with water. The reaction ranges from vigorous in case of Lithium to explosive in case of Cesium. • Melting and boiling points decrease down the group. • Density increases down the group. Lithium, Sodium and Potassium are less dense than water (density of water is 1 gcm-3) so they float over water. Elements Amani Hamdi Page 44 Group II(the alkaline earth metals) -Group II is called the alkaline earth metals. The elements of this group are: Beryllium, Magnesium, Calcium, Strontium, Barium, and the radioactive element Radium. Group II shows similar trends in reactivity to group 1, but are less reactive. -The reactivity increases going down the group. Beryllium (at the top) is the least reactive and Barium at the bottom is the most reactive. Group VII( Halogens) The most reactive non-metals are halogens in group VII of the periodic table They are all diatomic non-metals (for example, F2,Cl2,Br2,I2).. Physical properties of halogens • They have a similar strong smell. • Poor electric conductors. • Low melting and boiling points . Chemical properties of halogens • They have a valency of -1. • They react directly with metals to form halide salts. • They are all poisonous Trend of group VII • • • • The reactivity decreases down the group. The melting and boiling points increase down the group. Density increases down the group. The elements change from gases to liquid to solids as we go down the group. • The colours of the elements get darker as we go down the group. Elements Amani Hamdi Page 45 Halogen State Colour F2 Gas Pale yellow Cl2 Gas Yellow-green Br2 Liquid Red-brown I2 Solid Dark grey At2 Solid Black Transition elements -They are elements that lie between group II and group III beginning from the 4th period. -This block of elements includes many elements you are familiar with for example, Copper, Iron and Zinc. Physical properties of Transition elements -They are harder and stronger than metals of group I and II. -They have much higher densities than metals of group I and II. -They have high melting points (except for mercury which is liquid at room temperature). - A few of the metals are strongly magnetic (Iron,Cobalt & Nickel). Elements Amani Hamdi Page 46 Chemical properties of Transition elements. • Many of their compounds are coloured. • They have more than one valency (oxidation state); they form more than one type of ions for example, Iron can form Fe+2 and Fe+3 ions. • The transition elements and their compounds often make useful catalysts. • They are less reactive than the other metals Noble gases (group Zero) -The elements of this group are Helium, Neon, Argon, Krypton, Xenon and the radioactive element Radon. -They are non-metals. -They exist as monoatomic molecules (individual atoms). Physical properties of Noble gases -They are colorless gases -Their melting points and boiling points are extremely low. Chemical properties of Noble gases -They are very unreactive as they have completely filled outermost shell. Elements Amani Hamdi Page 47 The uses of Nobel gases depend on this inertness. • Helium is used in airships and balloons because it is both light and unreactive. • Argon is used to fill light bulb because it does not react with the filament even at high temperature. • Neon is used in advertising and lasers. Uses of some metals Uses Properties 1. Aluminum is used in cooking utensils Resists corrosion Good conductor of heat. It has a high tensile strength and light (low density) 2. Aluminum is used in the construction of aircraft. 5. Aluminum is used in making overhead power lines. Resists corrosion (as it is covered with non - porous protective oxide layer ) Good electric conductor and has low density. With a steel core To strengthen the cable 6. Cupper is used in making electric power cables. 7. Cupper is used in making cooking pans. Good electric conductor and ductile. 3. Aluminum is used in kitchen foil 4. Aluminum is used in Can 8. cupper is used in making of water pipes. 9. Iron is used in construction of bridges 10. Tin is used in coating tin cans Good heat conductor , has high melting point and resists corrosion. Cupper does not react with water and resists corrosion Hard. Not poisonous Elements Amani Hamdi Page 48 Recycling metals: Recycling a metal involves collecting used metal items and producing new metal from them . The steps usually needed are : Collecting and transporting the used items to a recycling centre Breaking up and sorting the different metals Removing impurities from the metals The metals can then be used to manufacture new metal items. Advantages of recycling metals: the advantage of recycling compared to producing metals from metal ore include : more economic - less energy is needed to produce a metal less damage to the environment - fewer quarries and mines less noise and less heavy traffic . saves valuable raw materials as reserves of metal ores will last longer Elements Amani Hamdi Page 49 Reactions of Metals with Air, Water and Dilute Acids We can get information on reactivity by investigating the following aspects of metal chemistry: • Ease of extraction. • Reactions with air and water. • Reactions with dilute acids. • Metal displacement reactions and Redox reactions. • Heat stability of metal compounds. The overall picture is summarized in what is known by the reactivity series of metals. The reactivity series of metals is an order of reactivity giving the most reactive metal first. Elements Amani Hamdi Page 50 Reaction of Metals with Air K Na Burn in air producing oxides. Ca Mg Al is a reactive metal that is covered with non-porous protective oxide layer. Al will burn only when the layer is removed. Zn Fe Sn Do not burn but hot metal glows and covers by oxide film Pb Cu Ag Au Do not react Pt Elements Amani Hamdi Page 51 Chemical equations to show the previous reactions Metal + Oxygen Metal Oxide Sodium + Oxygen Na + O2 Sodium Oxide (word equation) Na2O (symbolic equation) The previous equation is not balanced as the number of atoms of any element must be equal at both sides as matter is neither created nor destroyed. So you have to insert numbers in front of the chemicals to balance the equation: 4Na + O2 2Na2O Magnesium + Oxygen 2Mg + O2 Magnesium Oxide 2MgO Elements Amani Hamdi Page 52 Reactions of Metals and Water K Na Very reactive metals react with cold water producing Metal Hydroxide and Hydrogen. Ca 2K + 2H2O 2KOH + H2 2Na + 2H2O 2NaOH + H2 Ca + 2H2O Ca(OH)2+ H2 Mg Al React with steam producing Metal Oxide and Hydrogen. Al will react only when its non-porous oxide is removed. Zn Fe Mg + H2O MgO + H2 Sn Zn + H2O ZnO + H2 Pb Cu Ag Don’t react with water. Au Pt Elements Amani Hamdi Page 53 Note: Magnesium reacts very slowly with cold water, but reacts vigorously when heated with steam. Remember that as we go down, the group of alkali metals, the reactivity increases. Li Lithium reacts steadily with water; the metal does not melt and hydrogen does not ignite. Na Sodium reacts more strongly; the metal melts and Hydrogen Ignites. K Potassium reacts vigorously with water and Hydrogen ignites spontaneously. Potassium may explode dangerously. Increase s If you place a small piece of Sodium in water : 1. it floats and reacts very quickly. 2. The reaction is exothermic and the heat produced melts the Sodium. 3. The molten Sodium dart around the water surface 4. A yellow flame is seen. 5. You may see effervescence (fizz) as Hydrogen evolves. Elements Amani Hamdi Page 54 Displacement reactions of metals According to the reactivity series: More reactive metals can displace less reactive metals in their salt solutions. Mg(s) + CuSO4(aq) MgSO4(aq) + Cu(s) Magnesium is more reactive than Copper. Metals above Hydrogen in the chemical reactivity series react with dilute acids producing salt and Hydrogen. Zn(s) + 2HCl(aq) ZnCl2(aq)+ H2(g) Zinc is more reactive than Hydrogen Cu(s) + HCl(aq) No reaction Copper is less reactive than Hydrogen. Elements Amani Hamdi Page 55 Alloys An Alloy is a mixture of elements usually metals. Alloys are formed by mixing the molten metals together thoroughly and then allowing them to cool and form a solid. • Alloys are designed to have the properties useful for a particular purpose. • An alloy is stronger than the original individual metals. • Look at the fig. you will see that in case of pure metals, layers can slip easily over each other as all the layers are made of same sized atoms. In case of Alloys, the impurities (the different type of atoms) reduce the slippage because the atoms are of different size making the alloy harder than pure metals. Elements Amani Hamdi Page 56 The following table shows some important alloys, their composition, their properties and uses. Alloy Composition Copper 70% Brass Zinc 30% Copper 90% Bronze Tin Stainless steel Mild steel 10% Iron 74% Chromium 18% Nickel 8% Iron Carbon Properties • Harder than pure copper • Gold coloured • Harder than pur e copper • Harder than pure iron • Does not rust 99.7% • Stronger and 0.3% harder than pure iron Elements Uses • ornaments • ornaments • Cutlery • Surgery tools • Chemical plants • Car bodies • machinery Amani Hamdi Page 57 Rusting of iron and its prevention When a metal is attacked by air, water or other surrounding substances, it is said to corrode. In the case of iron and steel, the corrosion process is known as rusting. Rust is a red- brown powder consisting mainly of hydrated iron oxide (Fe2O3.2H2O). Water and oxygen are essential for iron to rust. Sea water increases the rate which iron objects rust. Rust prevention Iron can be prevented from rusting by means of many ways. 1 . By Barrier Methods : Isolating iron from oxygen and water a. Painting: Ships, Lorries, cars, bridges and many other iron and steel structures are painted to prevent rusting however, if the paint is scratched, the iron beneath it will start to rust. b. Oiling and greasing: The iron and steel in the moving parts of machinery are coated with grease to prevent them from rusting, but the protective film of grease must be renewed. Elements Amani Hamdi Page 58 c. Plastic coatings: These are used to form a protective layer on items such as refrigerators. d. Electroplating: An iron or steel can be electroplated with a layer of Chromium or Tin. A Tin can is made of steel coated in both sides with tin because tin is unreactive and non-toxic. If the protective layer is broken, then the steel beneath will begin to rust. e. By Galvanisation: Galvanisation is a process where iron is protected , when coated with a layer of zinc . This can be done by electroplating or dipping iron into zinc The zinc coat react with Oxygen producing Zinc oxide that react with the moisture producing Zinc Oxide that reacts with the moisture producing Zn(OH)2 , that react with CO2 producing ZnCO3 that protects iron by the barrier method . If the coating is damaged the is still protected by sacrificial protection. Elements Amani Hamdi Page 59 2. By Sacrificial protection: This is a method of rust prevention in which blocks of a reactive metal are attached to the iron surface. Zinc or Magnesium blocks are attached to oil rigs and to the halls of ships. The more reactive metal,Zinc or Magnesium will corrode in preference to iron. In sacrificial protection, the more reactive metal, Zinc or Magnesium loses electrons and oxides in preference to iron and these electrons flow to iron preventing iron from rusting forming. (Iron(III)Oxide). Elements Amani Hamdi Page 60 3. By Electrolytic protection (cathodic protection) Huge structures like oil-rig can be protected by using Iron as cathode (connected to the negative pole of the battery). Inert electrode is used as the anode (e.g. Titanium) and a power supply is required. The figure shows electrolytic protection of oil rig. Elements Amani Hamdi Page 61 Colours Elements Metals Compound Non-metals covalent Except: NO2 Except: Gold: Yellow F Yellow O2 ,H2 , N2 All are colourless Ionic C Black SYellow BWhite Br Red-brown Dark grey At Black Ionic compounds of group 1,2 ,3 are white Ionic compounds of coloured solution Note: Iodine is dark gray solid at room temperature .when sublimes, it gives purple fumes .In solutions it is Yellow-brown -Solutions of halides (ions of halogens) are colourless Colours Amani Hamdi Page 62 Colours of some transtion compounds Most Copper II compounds are blue Cu(OH)2 CuSO4.5H2O Cu(NO3)2 CuCl2 CuCO3 CuO Blue Blue - Green Green Black Most Iron II Compounds are green Fe(OH)2 FeO Most Iron III Compounds are red-brown Fe(OH)3 Fe2O3 Colours of some oxidizing agents -Potassium Chromate (VI) (K2Cr2O7) Orange Reduction -Potassium Manganate (VII) (KMnO4)Violet Reduction Green Colourless Colours of some hydrated salts Hydrated CoCl2.6H2O (Pink) Hydrated CuSO4.5H2O(Blue) Dehydration Dehydration Anhydrous CoCl2 (Blue) Used to test for water Anhydrous CuSO4 (White) Amani Hamdi Colours Page 63 Chemical reactions A chemical reaction is a change in which a new substance is formed. Generally the process is not easily reversed. During a chemical reaction, energy can be given out (exothermic reaction) or taken in(endothermic reaction). A physical change is the change in which the substances do not change identity. They can be easily returned to their original form by some physical processes. Examples of physical changes are: Change in state of matter. Melting, freezing, evaporation and condensation are physical change. Dissolving a substance in water. Grinding the substance such as salts or sugar. A chemical equation links together the names of substances that react (the reactants) with those of the new substances formed (products). Examples heat Magnesium + Oxygen Magnesium Oxide (Word equation) Reactants Products 2 Mg + O2 2 MgO (symbolic equation) This is a balanced equation, the number of each type of atoms are the same on both the reactant side and product side, two Magnesium atoms and two Oxygen atoms at each side. Amani Hamdi Chemical Reactions Page 64 Chemical Reactions Synthesis Combustion Displacement Single Displacement Decomposition Double Displacement Acid-Base Neutralization Ionic Precipitation Synthesis reactions This type occurs where two or more substances react together to from one product. Examples: Iron + Sulfur Iron(II)Sulfide Fe FeS + S Magnesium + Oxygen Magnesium Oxide 2Mg + O2 Hydrogen + Oxygen 2H2 + O2 2 MgO water 2 H2O Amani Hamdi Chemical Reactions Page 65 Note: Photosynthesis is a very essential process for life. Photosynthesis is a process by which plants make their food (glucose) by absorbing carbon dioxide from air, water from soil and sunlight trapped by chlorophyll. Chlorophyll acts as catalyst. It is a synthesis reaction. , It is an endothermic reaction. Carbon dioxide + water 6 CO2 sunlight Glucose + Oxygen sunlight + 6 H2O C 6 H12 O6 + 6O2 Displacement Reactions Double Displacement Reaction Simple Displacement Reaction Neutralization Reaction Precipetation Reaction Double displacement reaction: In double displacement reaction, two ionic compounds exchange their ions. A+B- (aq)+ C+D(aq ) AD + CB Amani Hamdi Chemical Reactions Page 66 A) Neutralization reaction: Neutralization is a chemical reaction between an acid and a base (or alkali) to produce salt and water only. Examples: HCl + NaOH Acid + alkali NaCl + H2O salt + water H2SO4+ CuO Acid + base CuSO4 + H2O salt + water Note: Metal Oxides and Metal Hydroxides are bases. Bases that dissolve in water become alkali. B) Precipitation Reaction: Precipitation reaction involves the formation of an insoluble product. Precipitation is the sudden formation of an insoluble salt when two solutions are mixed or when a gas is bubbled into a solution. -This type of reaction is used to prepare insoluble salts. Examples: Pb(NO3)2 (aq) + 2KI(aq) PbI2(s) + 2KNO3(aq) Insoluble salt Yellow ppt. -The downward arrow shows that PbI2is insoluble forming a precipitate. -The insoluble salt (PbI2) can be separated from the soluble salt (KNO3) by filtration. Amani Hamdi Chemical Reactions Page 67 -The lime water test for Carbon dioxide also depends on precipitation. When Carbon dioxide is passed through lime water Ca (OH)2 solution, an insoluble product CaCO3 is formed that causes turbidity. CO2(g)+ Ca(OH)2(aq) CaCO3(s) + H2O(l) Simple displacement reaction More reactive metal can displace a less reactive metal from its salts. Mg + CuSO4 MgSO4 + Cu (Magnesium is more reactive than cupper) Metals above hydrogen in the reactivity series are more reactive than hydrogen, so they can react with acids as the metal displaces hydrogen. Mg + 2HCl MgCl2 + H2 Halogens (elements of group VII) displace halides of less reactive halogens from their salts. F the most reactive halogen Cl Br I At least reactive halogen Amani Hamdi Chemical Reactions Page 68 Cl2 + 2KI Green gas Colourless soln. Cl2 + Green gas 2KBr 2 KCl + Brown liquid I2 colourless soln. Colourless soln. Br2 + 2 KCl + colourless soln. 2KI Colourless soln. Yellowish brown Br2 Brown liquid 2 KBr + colourless soln. I2 Yellowish brown Oxidation-Reduction Reactions (Redox Reactions) And Combustion Chemists’ ideas about oxidation and reduction have been expanded to a wide range. Oxidation is a chemical process in which a substance gains oxygen during a reaction. Reduction is a chemical process in which a substance loses oxygen during a chemical reaction. Example: heat Cupper (II) Oxide + Hydrogen Cupper + water reduction heat CuO + H 2 Cu + H2O Oxidation CuO is an oxidizing agent. H2 is reducing agent. Colour changes from black to reddish brown. Amani Hamdi Chemical Reactions Page 69 reduction 2Fe2O3+ 3C 4 Fe + 3 CO2 Fe2O3 is an oxidizing agent. Oxidation C is a reducing agent. The most common oxidizing agents are: Oxygen (or air), Hydrogen peroxide, Acidified Potassium Manganate (VII) and Acidified Potassium (VI)chromate. The most common reducing agents are: Hydrogen, Carbon, Carbon monoxide, and Sulfur dioxide - Aqueous potassium iodide KI (aq) Colorless solution of KI turns brown due to the formation of iodine 2I- (aq) I2 + 2eOxidation is the loss of electrons A new definition Remember this by : OIL Oxidation Is Losing e’ RIG Reduction Is gaining e’ Examples : Sodium + Chlorine 2Na Nao Clo + Cl2 Loses 1eGains 1e+ Na+ Cl- Sodium Chloride 2 Na+Cl(oxidation) (Reduction) During this reaction, the oxidation state of Sodium has increased by 1 from 0 to +1 and the oxidation state of Chlorine has decreased from 0 to –1 Amani Hamdi Chemical Reactions Page 70 Thus another definition for oxidation and reduction has to be included: Oxidation is the increase in oxidation state of an atom or an ion. Reduction is the decrease in oxidation state of an atom or an ion. e.g. reduction Zn0(s)+ Cu+2SO2-4(aq) Zn2+ SO2-4(aq)+ Cu0 oxidation Zn is oxidized. Cu2+ is reduced. The oxidation state of Zinc has increased by 2, from 0 to + 2. Meanwhile the oxidation state of Copper has decreased by 2, from + 2 to 0 e.g. Re reduction Cl02(g) + 2K1+I1-(aq) 2K1+Cl1-(aq) +I02(aq) oxidation -The oxidation state of iodide changes form –1 to 0. It has increased. Iodide has oxidized. -The oxidation state of Cl changes from 0 to –1. It has decreased. Chlorine is reduced to chloride ions. An oxidizing agent (Oxidant): a substance which will oxidize another in a redox reaction. An Reducing agent (Reductant ):a substance which will reduce another in a redox reaction. Amani Hamdi Chemical Reactions Page 71 Rules of Oxidation States Rule Example 1. The Oxidation number of any uncombined element is Zero H2 Zn O2 2. Many elements or ions have fixed oxidation number in compounds Group 1 elements are always +1 Group 2 elements are always +2 Fluorine is always -1 Hydrogen is +1 (except for in metal hydrides like NaH , where it is -1 ) Oxygen is -2 (except in peroxides , where it is -1 and F2O where it is +2 ) 3. The oxidation number of an element in monoatomic ion is always same as the charge Zn 2+ Oxidation number = +2 Fe 3+ oxidation = +3 Cl- oxidation number = -1 4. The sum of the oxidation number in a compound is zero NaCl Oxidation number Na = +1 Oxidation number Cl = -1 Sum oxidation number = 0 5. The sum of oxidation number in ion is equal to the charge on the ion SO4 2- Oxidation number of S = +6 Oxidation number 4 O atoms = 4x (-2) Sum oxidation number = -2 Amani Hamdi Chemical Reactions Page 72 Combustion Combustion: the reaction of a substance with Oxygen causing the release of energy. The reaction is exothermic and often involves a flame. Burning: combustion in which a flame is produced. -The combustion of natural gas is an important source of energy for homes and industry. Natural gas is mainly Methane. Its complete combustion produces Carbon dioxide and water vapour. Methane + Oxygen CH4 + 2 O2 carbon dioxide + water CO2 + 2H2O -Substances that undergo combustion readily and give out a large amount of energy are known as fuel. -Our bodies get their energy from food, in a process called respiration. Respiration: is the breaking down of food in body cells to release energy. Amani Hamdi Chemical Reactions Page 73 Energy and chemistry (Heat changes in chemical reactions) Energy cannot be created nor destroyed (law of conversation of energy). However energy can be changed from one form to another. Heat of reaction (heat change): It is the amount of energy taken in or given out, it is measured by kilojoule per mole (KJ/mol). 1 kilojoule = 1000 joule Heat energy is given the symbol ΔH (pronounced delta aitch) Energy changes occur in some chemical reactions or even in some physical processes such as when a solid dissolves in water. Chemical reactions in which heat is given out are described as exothermic and those in which heat is taken in are endothermic. Note: In an exothermic reaction, the temperature of the reaction mixture initially rises until the highest temperature is reached. When the reaction is completed, the temperature falls until it reaches the room temperature. In an endothermic reaction, the temperature of the reaction mixture falls until the lowest temperature is reached. When the reaction is completed the temperature of the reaction mixture rises until it reaches the room temperature. Energy And Chemistry Amani Hamdi Page 74 Exothermic reactions Definition Characteristics of the reaction Examples Endothermic reactions Reactions that give out heat energy to the surrounding. -Heat is liberated and is transferred from the chemicals to the surrounds. Reactions that absorb heat energy from the surrounding. -Heat is absorbed and is transferred from the surrounds to the reactants. -The temperature of the reaction mixture rises, the container feels hot. -The temperature of the reaction mixture falls, the container feels cold. -Heat of the reactants is more than heat of products. Combustion of fuel. Simple displacement reaction. Neutralization reaction. Reactions in cells. Respiration. Condensation. Freezing. Iron rusting. -Heat of the reactants is less than heat of products. Thermal decomposition. Electrolysis. The action of light on silver bromide in photographic film. Photosynthesis. Melting. Evaporation. If the reaction is exothermic, i.e. heat is given out; ΔH has a negative value (as the reactants lose energy). If the reaction is endothermic, i.e. heat is taken in; ΔH has a positive value (as the reactants gain energy). Energy And Chemistry Amani Hamdi Page 75 During a reaction, all bonds of the reactants are broken, this takes energy in, so Breaking chemical bonds takes energy from the surrounding. This is an endothermic process. New bonds are then formed to give products. Formation of bonds releases energy. Formation of bonds gives out energy to the surrounding. This is an exothermic process. Examples Burning methane in oxygen. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + + Bonds that are broken 4 C H bonds 2O O bonds Energy And Chemistry Bonds that are formed 2C O 4O H Amani Hamdi Page 76 If the energy absorbed in bond breaking is more than the energy Released in bond formation, so the reaction is endothermic. If the energy released in bond formation is more than the energy absorbed in bond breaking, so the reaction is exothermic. Bond energy: it is the energy required to break one mole of bonds. The overall change in energy for a reaction can be shown in an energy level diagram (or energy profile) Energy profile for an exothermic reaction: Burning Methane in Oxygen is an exothermic reaction. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) ΔH = negative value. Energy And Chemistry Amani Hamdi Page 77 Energy Profile for an endothermic reaction: The reaction between Nitrogen and Oxygen is an endothermic reaction. N2(g) + O2(g) 2NO(g) ΔH = positive value Energy And Chemistry Amani Hamdi Page 78 Energy Profile for an endothermic and exothermic reactions with activation energy Ea : Energy Profile for an endothermic and exothermic reactions with Catalyst: Ea : Activation energy without using a catalyst . E'a : Activation energy with using a catalyst . Energy And Chemistry Page 79 Amani Hamdi Worked example : Reaction between Hydrogen and Chlorine to produce Hydrogen Chloride. H2(g) + Cl2(g) 2HCl(g) If the bond energies of: Cl – Cl H–H H – Cl 242 KJ. 436 KJ. 431 KJ. Calculate the heat change and deduce if the reaction exothermic or endothermic. Solution: Energy taken in to break bonds For a mole of H2 molecule 436 KJ For a mole of Cl2 molecule 242 KJ Total energy in 678 KJ Energy given out from bond formation: For 2 moles of HCl 2 × 431 = 862 Energy in – energy out = + 678 KJ - 862 KJ = -184 KJ +ve sign as energy is gained -ve sign as energy is lost Generally: -ve sign indicates that the reaction is exothermic ΔH = + Energy in – Energy out So the reaction gives 184 KJ more energy than it takes in. Since it gives out energy, the reaction is exothermic. Energy And Chemistry Amani Hamdi Page 80 Acids and Bases An Acid: is a substance that dissolves in water producing hydrogen ions. Hydrogen ion is also known as a proton, so an acid can be defined as a proton donor. Acids present in animals and plant material are known as organic acids, e.g. ethanoic acid in vinegar and methanoic acid in antsand nettle stings. Acids in food make it sour, e.g. vinegar, lemon and grapefruit. Other acids are corrosive .They are dangerous on skin and some are able to attack metalwork and stonework. These powerful acids are called mineral acids, e.g. Hydrochloric acid and Sulfuric acid. Acids also can be classified into strong and weak acids. Strong acid: an acid that is completely ionized when dissolved in water producing high concentration of H+(aq) in solution. e.g. Hydrochloric acid. Weak acid: an acid that is partially dissociated in water producing low concentration of H+(aq) in solution. e.g. Ethanoic acid. Number of hydrogen ions that can be donated by one molecule of acid is called basicity. Acid and Alkalis Amani Hamdi Page 81 Common acids Acid Formula Ions present Basicity Hydrochloric acid HCl H+ , Cl( chloride) Monobasic Nitric acid HNO3 H+, NO-3 (nitrate) Monobasic H2SO4 2H+, SO4-2 (sulfate) Dibasic 3H+, PO4-3 (phosphate) Tribasic Sulfuric acid Phosphoric acid H3PO4 Ethanoic acid CH3COOH H+ , Acid and Alkalis CH3COO(ethanoate) Monobasic Amani Hamdi Page 82 Reactions of Acids Reaction Metal + acid Salt + hydrogen Equation Magnesium+ Hydrochloric acid Magnesium chloride + Hydrogen Mg(s) + 2HCl(aq) MgCl2(aq) + Base + acid Sodium hydroxide + Hydrochloric acid Salt + water NaOH(aq) + HCl(aq) (neutralization reaction) Carbonate salt + acid Calcium carbonate + Sulfuric acid salt + water + carbon dioxide CaCO3(s) + H2SO4(aq) H2(g) Observation Bubbles of Colourless gas. Lighted splint burns with pop sound. Sodium Chloride + water NaCl(aq) + H2O(l) Calcium sulfate + water + Carbon dioxide CaSO4(s) + H2O(l) + CO2(g) Effervescence of a gas that Turns lime water milky. Note: if you are asked to suggest a metal with an acid, do not go with very reactive metals such as Ca. Acid and Alkalis Amani Hamdi Page 83 Bases Most metal oxides and metal hydroxides are bases. A base is a substance that neutralizes an acid to form a salt and water only. A base is a proton acceptor. (Remember that an acid is a proton donor. A proton is H+) NH3 + H2O NH4+ + OHBase (proton acceptor) Acid (proton donor) In this equation you will find that H+ is transferred from H2O to NH3, So NH3 here acts as a base because it is a proton acceptor and water acts as an acid because it is a proton donor. HCl + H2O H3O+ + ClAcid (proton donor) Base (proton acceptor) Most bases are insoluble in water; those that dissolve in water are called alkalis. The common alkalis are NaOH, KOH, Ca(OH)2 and NH4OH An alkali is a soluble base which produces OH-(aq) ion in water. Alkalis feel soapy to skin because they convert the oil in your skin into soap. An alkali turns red litmus paper blue. Alkalis are either strong or weak. Strong alkali is an alkali that ionizes completely when dissolved in water producing high concentration of OH-(aq) in the solution. For example Sodium Hydroxide Weak alkali is an alkali that dissociates partially giving low concentration of OH-(aq) ions in the solution. For example: Ammonium Hydroxide. Acid and Alkalis Amani Hamdi Page 84 Some common alkalis and bases Type Alkali Bases Name Formula Ions Strength Sodium hydroxide NaOH Na+, OH- Strong Potassium hydroxide KOH K+, OH- Strong Ammonium hydroxide NH4OH NH4+, OH- Weak Calcium Oxide CaO Don’t ionize strong Magnesium Oxide MgO Acid and Alkalis Amani Hamdi Page 85 Reactions of bases and alkalis Reaction Acid + base salt + water (neutralization) Alkali + ammonium salt heating Salt + water + ammonia Alkali + fat salt of fatty acids ( Soap)+ glycerol Equation Sulfuric acid + Magnesium Oxide H2SO4 + MgO Sodium Hydroxide + Ammonium Chloride NaOH(aq) + NH4Cl(S) NaOH(aq) + Oil Observation Magnesium Sulfate + water MgSO 4 + H2O The released Sodium Chloride+ water+Ammonia ammonia can be tested using NaCl(aq) + H2O(L) + NH3(g) damp red litmus paper, it turns blue. salt of fatty acids(soap) + glycerol Acid and Alkalis Amani Hamdi Page 86 Oxides when an element combines with oxygen, an oxide is formed. Types of oxides Metal oxides Basic Oxides Non-metal oxides Acidic Oxides Amphoteric Neutral Oxide Amani Hamdi Acid and Alkalis Page 87 Basic Amphoteric Examples Most metal Oxides (or hydroxides) E.g: CaO , MgO ,FeO , Fe2O3, CuO ,etc. ZnO Al2O3 Most nonmetal Oxides e.g: CO2, SO2,NO2, etc. Acidic Neutral H2O NO CO Type of oxide Properties Ionic compound. Those that dissolve in water produce alkalis. React with acids producing salt and water Only (Neutralization reaction) Ionic compounds. React with both acids alkalis. Covalent Compounds. Dissolve in water Producing acids. React with alkali producing salt and water only ( Neutralization reaction) Do not react with acids. Covalent compounds. They neither react with acids nor alkalis. Amani Hamdi Acid and Alkalis Page 88 Indicators An indicator is used to detect whether a substance is acidic or basic. An indicator is a substance that changes colour when added to an acid or to an alkali solution. Some common indicators and their colour change indicator Litmus phenolphthalein Methyl orange Thymophthalein Acidic medium Red Colorless Red Colorless Neutral medium Purple Colorless Orange Colorless Alkaline medium Blue Pink Yellow Blue Universal Indicator(full-range indicator) It gives range of colours (a spectrum) depending on the strength of the acid or the alkali. Note: solutions of the same acid but of different concentrations give different colours. Acid and Alkalis Amani Hamdi Page 89 The pH scale It is a scale running from 0 to 14. Used for expressing the strength of acids and alkalis. Acids have pH less than 7. The more acidic a solution, the lower the pH. Neutral substances, such as pure water have a pH =7. Alkalis have a pH greater than 7. The more alkaline a solution, the higher its pH. The pH of a solution can be measured using universal indicator (solution or paper). More accurate method to measure the pH is to use a pH meter. How the colour of universal indicator changes in solutions of different pH values Acid and Alkalis Amani Hamdi Page 90 Salt preparation -Salts are ionic compounds made by neutralization of an acid with a base (or alkali). -To prepare a salt, you have to see if the salt is soluble or insoluble. -The following diagram shows the solubility of various types of salts. Preparing soluble salts Salt Preparation Amani Hamdi Page 91 Note Salts including potassium , sodium & Ammonium ions are soluble in water. All nitrate salts are soluble in water . All sulfate salts are soluble except pb , Ba & (Ca). All chloride salts are soluble except Pb and Ag . All carbonate salts are insoluble excepts Na , K and NH 4 . All Oxide and hydroxide salts are insoluble except Na , K , NH4 and (Ca) Hydrated substance : substance that is chemically combined with water Anhydrous substance : substance containing no water Water of Crystallisation : water molecules present in hydrated Crystals CuSO4 .5 H 2O and CoCl2 . 6H 2O Salt Preparation Amani Hamdi Page 92 Soluble salts can be prepared from their parent acids sing any of the characteristic reactions of acids mentioned before: Metal + Acid salt + Hydrogen Base + Acid salt + water Carbonate salt + Acid salt + water + Carbon dioxide There are two methods to prepare soluble salts Method A: This method is used to prepare a soluble salt from insoluble metal, metal Oxide (base) or Carbonate salt. Follow the following steps 1- Add excess (more than enough) of the solid (metal, metal oxide or metal carbonate) to the acid. Excess solid is added to make sure that all the acid is used up. The excess is seen at the bottom. 2- The excess solid is removed by filtration. 3- The filtrate is gently evaporated till point of crystallization and get a saturated solution . 4- The concentrated solution obtained is left to cool to obtain crystals. The crystals are filtered off. 5- Crystals are dried carefully between two filter papers. saturated solution : solution in which no more solute is dissolved at certain temperature . Salt Preparation Amani Hamdi Page 93 Remember to choose the suitable parent acid Salt of Chloride Sulfate Nitrate Parent acid Hydrochloric acid Sulfuric acid Nitric acid Salt Preparation Amani Hamdi Page 94 Steps of preparing crystals of MgSO4.7H2O 1- Add excess Mg MgO MgCO3 to the parent acid (Sulfuric acid) In case of using the base MgO, you have to heat the acid to start the reaction. No need to heat the acid if you use the metal or the carbonate salt because they are readily to react with acid at room temperature. 2- Filter to remove the excess Mg, MgO, MgCO3. 3- Half evaporate the filtrate till point of crystallization. 4- Leave to cool to get the crystals. Filter off the crystals. 5- Dry the crystals gently between two filter papers. Generally The solid added Is there a need to Reason to the acid heat the acid Readily to react at Metal No room temperature To start the Metal Oxide Yes reaction Metal Carbonate No What shows that the reaction stops Bubbles stop releasing. Excess metal is seen at the bottom. Excess metal oxide is seen at the bottom. Readily to Effervescence stops react at releasing. room The excess metal temperature carbonate is seen at the bottom. Salt Preparation Amani Hamdi Page 95 Method B : Acid + Alkali by Titration This method is used to prepare a soluble salt from a soluble base (alkali) or carbonate. Since both reactants are colourless and the alkali or the carbonate are soluble, an indicator must be used. Follow the following steps: 1- The acid is poured in a burette till its 0.0cm3. A known volume of an alkali is delivered into a conical flask using a pipette. Drops of an indicator are added for example Methyl Orange 2- The acid is run from the burette into the flask until the indicator just changes colour. The volume of the acid ran into the flask is recorded. The experiment is then repeated without using the indicator using the known volume of the alkali used as well as the recorded acid volume. Note: instead of repeating the experiment, activated charcoal is added to remove the indicator, and then the charcoal is filtered off. 3- The salt solution is evaporated in the evaporating dish till point of crystallization. Cool to form crystals. Filter off the crystals from the remaining solution, wash with distilled water. Dry the crystals gently between two filter papers. Note : universal indicator is not used in titration as an indicator as it gives a range of colors ,so the end point is not sharp . Salt Preparation Amani Hamdi Page 96 Steps of preparing crystals of Sodium Chloride from the soluble base Sodium Hydroxide: 1- Fill a burette till its 0.0cm3 with Hydrochloric acid. 2- Deliver a known volume of Sodium Hydroxide in a conical flask using a pipette. 3- Add drops of Methyl Orange to the flask. Methyl Orange is yellow in alkaline media and orange in neutral ones. 4- Add the acid from the burette with regular shaking until the colour changes from yellow to orange 5- Note the volume of acid used is for neutralization. 6- Repeat the experiment without the indicator using the same volume of acid recorded as well as the alkali. 7- Evaporate the resulting solution in an evaporating dish till point of crystallization. 8- Leave the concentrated solution to cool for the crystals to form. 9- Filter off the crystals from the remaining solution. Wash the crystals with distilled water. 10- Dry the crystals carefully between two filter papers. Salt Preparation Amani Hamdi Page 97 Note : Do not dry the crystals in an oven or by heating to avoid dehydration ( losing water of crystallization) The water of crystallization give the crystals their shape. In some cases it also gives them their colour for example: CuSO4.5H2O heat CuSO4 -5H2O (blue hydrated salt) (white anhydrous salt) When hydrated salts are heated: Their water of crystallization driven off as steam. The crystals lose their shape and become powder. Sometimes the colour changes. If the water is added again: The white anhydrous Copper(II)Sulfate changed to blue.(This can be used as test for water). Heat is given off. Salt Preparation Amani Hamdi Page 98 Preparation of insoluble salts Insoluble salts are made by precipitation. Examples: AgCl,BaSO4,PbI2 Steps of preparing Silver Chloride: Before writing the steps let us select two reagents (chemicals used) to prepare Silver Chloride. Silver Chloride Ag+ ClSelect a soluble ionic compound containing Cl-, like NaCl and make an aqueous solution of NaCl. Select a soluble ionic compound Containing Ag+, like AgNO3And make an aqueous solution of AgNO3. Steps : 1. Mix the two solutions of Silver Nitrate and Sodium Chloride, a precipitate is formed (AgCl). 2. Filter off the precipitate to get it as residue. 3. Wash the residue with distilled water. 4. Dry the residue in an oven. Salt Preparation Amani Hamdi Page 99 The equation for this reaction is: Silver Nitrate + Sodium Chloride Sodium Nitrate + Silverchloride AgNO3(aq) + NaCl(aq) NaNO3(aq) + AgCl(s) The state symbol here is very important, it shows you the precipitation The equation can be simplified to show only those ions that take part in the reaction and their products. The resulting equation is called ionic equation. How to write an ionic equation: 1- Write the equation with its state symbols AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) 2- Write a full ionic equation by separating the ions from the aqueous compounds, as they are dissociated in the solutions. Ag+(aq) + NO-3(aq) + Na+(aq) + Cl-(aq) Na+(aq) + NO-3(aq) + AgCl(s) AgCl is insoluble so it will not dissociate into ions. 3- Cross the ions that remain unchanged, they are called spectator ions. Ag+(aq) + NO-3(aq) + Na+(aq) + Cl-(aq) Na+(aq) + NO-3(aq) + AgCl(s) 4- Write the net ionic equations without the spectator ions. Ag+(aq) +Cl-(aq) AgCl(s) Salt Preparation Amani Hamdi Page 100 State symbol Symbol S l g aq Meaning solid liquid gas Aqueous solution i.e. dissolved in water. Write the net ionic equation of the following reaction Equation: Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s) Full ionic equation: Pb+2(aq) + 2NO3-(aq) + 2K+(aq) +2I-(aq) 2K+(aq) + 2NO3- + PbI2(s) Net ionic equation: Pb+2(aq)+2I-(aq) PbI2(s) Salt Preparation Amani Hamdi Page 101 Mixture separation and purification -A mixture must be made of at least two substances, which may be solid, liquid or gas. There are many ways in which the three matter states can mix. If the states (solid, liquid or gases) are completely mixed to become one phase, this mixture is called solution. Example, the solid salt dissolves in liquid water to produce a liquid mixture. -The solid that dissolves in the liquid is the solute. -The liquid in which the solid dissolves is called the solvent. -In other types of mixture, the states remain separate. An example of this type of mixture is a suspension of fine particles of a solid in a liquid, such that we get in a precipitation reaction. The most useful separation method for a particular mixture depends on : The type of mixture The substance in the mixture we are interested in. Types of mixtures Mixture of solids Mixture of solid and liquid Mixture of liquids Mixture of gases Mixture separation and purification Amani Hamdi Page 102 Mixture of solids A) Two solids one is soluble(A), the other is insoluble(B) Steps of separation The mixture is delivered into a beaker. Add excess water to dissolve (A) and stir. Filter, the residue is B and the filtrate is A. Crystalize the filtrate to get solid crystals of A. Note: Complete evaporation will produce powder of A. B) Separating two solids one is magnetic, the other is non-magnetic. E.g. A mixture of Iron (magnetic) and Sulfur (non-magnetic). Use a magnet to attract and separate iron from Sulfur. C) Separating two solids, one can sublime. E.g. A mixture of ammonium chloride and sodium chloride. By heating, ammonium chloride sublimes leaving sodium chloride. Mixture separation and purification Amani Hamdi Page 103 Mixture of a solid and a liquid A) The solid is insoluble in the liquid. E.g. Calcium Carbonate and water. The mixture can be separated by filtration. The solid is the residue and the liquid is the filtrate. Decantation Also can be used to separate a liquid from an insoluble salt. Decantation is The process of removing a liquid from a solid by pouring the liquid carefully . B) The solid dissolved in the liquid. -If you want to get the solid only, crystallize the solution. E.g. CuSO4.5H2O crystals from aqueous CuSO4. -If you want to get both liquid and solid, carry out simple distillation. Simple distillation: is the process of boiling a liquid and then condensing the vapour produced back into a liquid. A condenser is used to condense the vapour into liquid. The liquid is distillate. The process is also used to purify liquids. Mixture separation and purification Amani Hamdi Page 104 Mixtures of Liquids Miscible Liquids Immiscible Liquids A) Separation of miscible liquids: This separation is based on the fact that the liquids have different boiling points. The process used is known as fractional distillation. -The most volatile liquid (with lower boiling point) in the mixture distils over first and the least volatile liquid (with higher boiling point) distils over last. E.g. a mixture of ethanol and water. Ethanol boils at 78⁰C whereas water boils at 100⁰C. Ethanol distils first and collected then water. -Liquefied air is separated by fractional distillation. -Petroleum also is separated by the same technique. -The fractionating column has glass beads in the column to provide a large surface area for condensation. Mixture separation and purification Amani Hamdi Page 105 B) Separation of immisble liquids The process of separation depends on the different densities of liquids. E.g. a mixture of water and oil. Oil has less density than water. Separating funnel is used to separate such mixtures. The lower (denser) layer is tapped off at the bottom. This type of separation is used in industry. E.g. in the blast furnace, the molten slag forms a separate layer on top of molten iron. The two layers can then be tapped off separately. Mixture of Gases Mixture of gases can be separated either by: A) Diffusion: This way depends on the difference in Mr of the molecules of the gases. The one that has less Mr diffuses faster. B) Fractional distillation of Liquefied mixture of gase:s This way depends on the difference in boiling points. E.g. Components of air can be separated by fractional distillation of liquefied air. Mixture separation and purification Amani Hamdi Page 106 Separating coloured solids dissolved in a solvent by chromatography E.g. Separation of the colours present in a dye in ink. To separate using chromatography follow the following steps: 1. Get a chromatography paper. 2. Draw a line with a pencil (insoluble in solvents) called base line. Do not use ink. 3. Put a drop of the sample on the base line using a glass rod. 4. Put the paper in a suitable solvent .The level of solvent must start below the sample to prevent the solubility of the spot in the solvent. 5. Leave the chromatogram for several hours. Spots of several colours are separated. base line A drop of dye solvent Solvent front 3 separated sports Mixture separation and purification Amani Hamdi Page 107 Notes: Substances separate according to their solubility in the solvent. The dyes are carried by the solvent and begin to separate. The substance that is most soluble moves the fastest up the paper. An insoluble substance will remain at the origin. The run stops just before the solvent front (the level reached by the solvent). If the separated substances are colourless, the paper is treated with locating agent that reacts with the samples to produce coloured spots e.g Amino acids are colourless, a locating agent is used. Note : Molecules of amino acids can also be viewed under ultra violet light. The purity and identity of substances: The distance moved by a particular spot is measured and related to the position of the solvent front. The ratio of these distances is called rate of flow (Rf )value. This value is used to identify the substance: Rf = Distance moved by the sample Distance moved by the solvent Paper chromatography is one test that can be used to check the Purity of a substance. If the sample is pure, it should give one spot by chromatography. The identity of substances can also be checked by comparing its Rf value to that of a pure sample of the same substance. Mixture separation and purification Amani Hamdi Page 108 Practical work Safety precautions Never run around the laboratory. Never drink or eat in the laboratory. Wear safety goggles. Wear white coat. Tie hair. Make sure where the fire extinguishers are and how to use them. Never taste any chemicals in the laboratory. Never handle chemicals with your fingers ( use clean dry spatula ). Use a fume cupboard in lab or well ventilated lab if harmful gases are produced like halogens. Don’t heat alcohols directly on the Bunsen burner but use water bath. Chemical warnings and signs: Explosive These substances may explode if ignited in air or exposed to heat. A sudden shock or friction may also start an explosion. Particular care should be taken in warm climates. Oxidizing These substances produce much heat as they react with other materials. They can create a fire risk. They will react with organic substances such as wood. Practical work Amani Hamdi Page 109 Highly flammable These are solids, liquids or gases that may easily catch fire in a laboratory under normal conditions. Poisonous These substances are a serious risk to health. Some can cause death. These chemicals can affect you if you swallow them, inhale them, or absorbed through the skin. They should be stored in a locked cupboard. Harmful These chemicals are less of a health risk than poisonous substances, but they must be handled with care. Corrosive These chemicals destroy living tissues, including eyes and skin. If they are spilt on your skin, you should wash the exposed area with plenty of water. Irritant These substances are not corrosive but they can cause reddening of the skin. The effect may be immediate or it may only be observed after prolonged, or repeated contact with the chemical. Radioactive The radioactive chemicals used in schools have low activity. They are normally only used by teachers, for demonstrations. They should be treated in the same way as poisonous substances. Practical work Amani Hamdi Page 110 Lab equipment: Test tube holds liquids/solids Beaker holds liquids Funnel for pouring/filtering Boiling tube for heating liquids Triangle support for crucible Beehive shelf for collecting gas over water Dropper for measuring small volumes of liquid Burette for measuring large exact volume of liquid Crucible for heating solids to high temperature Gas jar to collect gases Practical work Amani Hamdi Page 111 Gauze spreads heat under a flask Measuring cylinder for measuring liquids Tripod supports beaker, crucible etc… over flame Evaporating dish for evaporating solutions Gas syringe for measuring volume of gases Round bottom flask for heating liquids Pipette for delivering exact volume of liquid Conical flask to hold liquids Tap funnel for separating immiscible liquids Tong to hold objects Practical work Amani Hamdi Page 112 Condenser for condensing vapour Spatula for transferring solids Mortar and pistil for grinding large solid substances Stand supports flask, tube etc… in complex apparatus Improvements for more accurate experimental results: Use Styrofoam (polystyrene) cup or insulated cups in experiments that are involved in heat change. Burette is used to measure the volume of solution added accurately rather than measuring cylinder. Pipette delivers a fixed volume accurately rather than measuring cylinder. Practical work Amani Hamdi Page 113 Methods of collecting gases 1) Downward delivery is used for gases that areser den than air such as: Chlorine, Hydrogenbo Chloride, Car dioxide, Sulfur dioxide andn Nitrogen dioxide. 2) Upward delivery is used for gases that are less dense than air such as Ammonia and Hydrogen. 3) Collection over water is used for gases that are not very soluble in water such as: Hydrogen, Oxygen and Nitrogen. 4) Collection in a gas syringe is useful when volume of gas needs to be measured. Practical work Amani Hamdi Page 114 Danger of “back suction” The problem arises when heating is stopped before delivery tube is removed from water. The reduced pressure in the reaction tube as it cools; results in the cold water can be sucked back into the hot boiling tube. The tube will crack and an explosion may occur. “Sucking back” can be prevented by making sure that the delivery tube is removed first before heating is stopped. Practical work Amani Hamdi Page 115 Functions of ceramic wool 1) To hold liquids. Delivery tube 2) Act as a separator. 3) Act as filter to prevent spreading of the reactants contaminating the rest of the apparatus. Practical work Amani Hamdi Page 116 Methods of drying gases: 1) Conc. Sulfuric acid is used to dry all gases except ammonia. 2) Anhydrous calcium chloride is used for all gases except ammonia, which forms a complex with calcium chloride. 3) Calcium oxide is used to dry ammonia and neutral gases. Practical work Amani Hamdi Page 117
