JK-TK/Newdata huggingface dataset:
They are all non metals. They include: Element Symbol Atomic number Electronicc configuration Charge of ion Valency State at Room Temperature Fluorine Chlorine Bromine Iodine Astatine F Cl Br I At 9 17 35 53 85 2:7 2:8:7 2:8:18:7 2:8:18:18:7 2:8:18:32:18:7 F- Cl- Br- I- At- 1 1 1 1 1 Pale yellow gas Pale green gas Red liquid Grey Solid Radioactive All halogen atoms have seven electrons in the outer energy level. They acquire/gain one electron in the outer energy level to be stable. They therefore are therefore monovalent .They exist in oxidation state X- The number of energy levels increases down the group from Fluorine to Astatine. The more the number of energy levels the bigger/larger the atomic size. e.g. The atomic size/radius of Chlorine is bigger/larger than that of Fluorine because Chlorine has more/3 energy levels than Fluorine (2 energy levels). 17 Atomic radius and ionic radius of Halogens increase down the group as the number of energy levels increases. The atomic radius of Halogens is smaller than the ionic radius. This is because they react by gaining/acquiring extra one electron in the outer energy level. The effective nuclear attraction on the more/extra electrons decreases. The incoming extra electron is also repelled causing the outer energy level to expand to reduce the repulsion and accommodate more electrons. Table showing the atomic and ionic radius of four Halogens Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Fluorine F 9 0.064 0.136 Chlorine Cl 17 0.099 0.181 Bromine Br 35 0.114 0.195 Iodine I 53 0.133 0.216 The atomic radius of Chlorine is 0.099nM .The ionic radius of Cl- is 0.181nM. This is because Chlorine atom/molecule reacts by gaining/acquiring extra one electrons. The more/extra electrons/energy level experience less effective nuclear attraction /pull towards the nucleus .The outer enegy level expand/increase to reduce the repulsion of the existing and incoming gained /acquired electrons.
Table showing the atomic and ionic radius of four Halogens Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Fluorine F 9 0.064 0.136 Chlorine Cl 17 0.099 0.181 Bromine Br 35 0.114 0.195 Iodine I 53 0.133 0.216 The atomic radius of Chlorine is 0.099nM .The ionic radius of Cl- is 0.181nM. This is because Chlorine atom/molecule reacts by gaining/acquiring extra one electrons. The more/extra electrons/energy level experience less effective nuclear attraction /pull towards the nucleus .The outer enegy level expand/increase to reduce the repulsion of the existing and incoming gained /acquired electrons. Electronegativity The ease of gaining/acquiring extra electrons is called electronegativity. All halogens are electronegative. Electronegativity decreases as atomic radius increase. This is because the effective nuclear attraction on outer electrons decreases with increase in atomic radius. The outer electrons experience less nuclear attraction and thus ease of gaining/acquiring extra electrons decrease. It is measured using Pauling’s scale. Where Fluorine with Pauling scale 4.0 is the most electronegative element and thus the highest tendency to acquire/gain extra electron. Table showing the electronegativity of the halogens. Halogen F Cl Br I At Electronegativity (Pauling scale) 4.0 3.0 2.8 2.5 2.2
18 The electronegativity of the halogens decrease down the group from fluorine to Astatine. This is because atomic radius increases down the group and thus decrease electron – attracting power down the group from fluorine to astatine. Fluorine is the most electronegative element in the periodic table because it has the small atomic radius. Electron affinity The minimum amount of energy required to gain/acquire an extra electron by an atom of element in its gaseous state is called 1st electron affinity. The SI unit of electron affinity is kilojoules per mole/kJmole-1 . Electron affinity depend on atomic radius. The higher the atomic radius, the less effective the nuclear attraction on outer energy level electrons and thus the lower the electron affinity.
The SI unit of electron affinity is kilojoules per mole/kJmole-1 . Electron affinity depend on atomic radius. The higher the atomic radius, the less effective the nuclear attraction on outer energy level electrons and thus the lower the electron affinity. For halogens the 1st electron affinity decrease down the group as the atomic radius increase and the effective nuclear attraction on outer energy level electrons decrease. Due to its small size/atomic radius Fluorine shows exceptionally low electron affinity. This is because a lot of energy is required to overcome the high repulsion of the existing and incoming electrons. Table showing the election affinity of halogens for the process X + e -> X- The higher the electron affinity the more stable theion.i.e Cl- is a more stable ion than Br- because it has a more negative / exothermic electron affinity than Br- Electron affinity is different from: (i) Ionization energy. Ionization energy is the energy required to lose/donate an electron in an atom of an element in its gaseous state while electron affinity is the energy required to gain/acquire extra electron by an atom of an element in its gaseous state. (ii) Electronegativity. -Electron affinity is the energy required to gain an electron in an atom of an element in gaseous state. It involves the process: X(g) + e -> X-(g) Electronegativity is the ease/tendency of gaining/ acquiring electrons by an element during chemical reactions. Halogen F Cl Br I Electron affinity kJmole-1 -333 -364 -342 -295
19 It does not involve use of energy but theoretical arbitrary Pauling’ scale of measurements. Physical properties State at room temperature Fluorine and Chlorine are gases, Bromine is a liquid and Iodine is a solid. Astatine is radioactive . All halogens exist as diatomic molecules bonded by strong covalent bond. Each molecule is joined to the other by weak intermolecular forces/ Van-der-waals forces. Melting/Boiling point The strength of intermolecular/Van-der-waals forces of attraction increase with increase in molecular size/atomic radius. Iodine has therefore the largest atomic radius and thus strongest intermolecular forces to make it a solid. Iodine sublimes when heated to form (caution: highly toxic/poisonous) purple vapour.
Melting/Boiling point The strength of intermolecular/Van-der-waals forces of attraction increase with increase in molecular size/atomic radius. Iodine has therefore the largest atomic radius and thus strongest intermolecular forces to make it a solid. Iodine sublimes when heated to form (caution: highly toxic/poisonous) purple vapour. This is because Iodine molecules are held together by weak van-derwaals/intermolecular forces which require little heat energy to break. Electrical conductivity All Halogens are poor conductors of electricity because they have no free delocalized electrons. Solubility in polar and non-polar solvents All halogens are soluble in water(polar solvent). When a boiling tube containing either chlorine gas or bromine vapour is separately inverted in a beaker containing distilled water and tetrachloromethane (non-polar solvent), the level of solution in boiling tube rises in both water and tetrachloromethane. This is because halogen are soluble in both polar and non-polar solvents. Solubility of halogens in water/polar solvents decrease down the group. Solubility of halogens in non-polar solvent increase down the group. The level of water in chlorine is higher than in bromine and the level of tetrachloromethane in chlorine is lower than in bromine. Caution: Tetrachloromethane , Bromine vapour and Chlorine gas are all highly toxic/poisonous. 20 Table showing the physical properties of Halogens Halogen Formula of molecule Electrical conductivity Solubility in water Melting point(oC) Boiling point(oC) Fluorine F2 Poor Insoluble/soluble in tetrachloromethane -238 -188 Chlorine Cl2 Poor Insoluble/soluble in tetrachloromethane -101 -35 Bromine Br2 Poor Insoluble/soluble in tetrachloromethane 7 59 Iodine I2 Poor Insoluble/soluble in tetrachloromethane 114 sublimes Chemical properties (i)Displacement Experiment Place separately in test tubes about 5cm3 of sodium chloride, Sodium bromide and Sodium iodide solutions.
The level of water in chlorine is higher than in bromine and the level of tetrachloromethane in chlorine is lower than in bromine. Caution: Tetrachloromethane , Bromine vapour and Chlorine gas are all highly toxic/poisonous. 20 Table showing the physical properties of Halogens Halogen Formula of molecule Electrical conductivity Solubility in water Melting point(oC) Boiling point(oC) Fluorine F2 Poor Insoluble/soluble in tetrachloromethane -238 -188 Chlorine Cl2 Poor Insoluble/soluble in tetrachloromethane -101 -35 Bromine Br2 Poor Insoluble/soluble in tetrachloromethane 7 59 Iodine I2 Poor Insoluble/soluble in tetrachloromethane 114 sublimes Chemical properties (i)Displacement Experiment Place separately in test tubes about 5cm3 of sodium chloride, Sodium bromide and Sodium iodide solutions. Add 5 drops of chlorine water to each test tube: Repeat with 5 drops of bromine water instead of chlorine water Observation Using Chlorine water -Yellow colour of chlorine water fades in all test tubes except with sodium chloride. -Coloured Solution formed. Using Bromine water Yellow colour of bromine water fades in test tubes containing sodium iodide. -Coloured Solution formed. Explanation The halogens displace each other from their solution. The more electronegative displace the less electronegative from their solution. Chlorine is more electronegative than bromine and iodine. On adding chlorine water, bromine and Iodine are displaced from their solutions by chlorine. Bromine is more electronegative than iodide but less 6than chlorine. 21 On adding Bromine water, iodine is displaced from its solution but not chlorine.
On adding chlorine water, bromine and Iodine are displaced from their solutions by chlorine. Bromine is more electronegative than iodide but less 6than chlorine. 21 On adding Bromine water, iodine is displaced from its solution but not chlorine. Table showing the displacement of the halogens (V) means there is displacement (x ) means there is no displacement Chemical /ionic equations With Fluorine F2(g) + 2NaCl-(aq) -> 2NaF(aq) + Cl2(aq) F2(g) + 2Cl-(aq) -> 2F-(aq) + Cl2(aq) F2(g) + 2NaBr-(aq) -> 2NaF(aq) + Br2(aq) F2(g) + 2Br-(aq) -> 2F-(aq) + Br2(aq) F2(g) + 2NaI-(aq) -> 2NaF(aq) + I2(aq) F2(g) + 2I-(aq) -> 2F-(aq) + I2(aq) With chlorine Cl2(g) + 2NaCl-(aq) -> 2NaCl(aq) + Br2(aq) Cl2(g) + 2Br-(aq) -> 2Cl-(aq) + Br2(aq) Cl2(g) + 2NaI-(aq) -> 2NaCl(aq) + I2(aq) Cl2(g) + 2I-(aq) -> 2Cl-(aq) + I2(aq) With Bromine Br2(g) + 2NaI-(aq) -> 2NaBr(aq) + I2(aq) Br2(g) + 2I-(aq) -> 2Br-(aq) + I2(aq) Halogen ion in solution Halogen F- Cl- Br- I- F2 X Cl2 X X Br2 X X X I2 X X X X
22 Uses of halogens (i) Florine – manufacture of P.T.F.E (Poly tetra fluoroethene) synthetic fiber. - Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass.
- Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass. (ii) Bromine - Silver bromide is used to make light sensitive photographic paper/films. (iii) Iodide – Iodine dissolved in alcohol is used as medicine to kill bacteria in skin cuts. It is called tincture of iodine. The table below to show some compounds of halogens. (i) Below is the table showing the bond energy of four halogens. Bond Bond energy k J mole-1 Cl-Cl 242 Br-Br 193 I-I 151 I. What do you understand by the term “bond energy” Bond energy is the energy required to break/ form one mole of chemical bond II. Explain the trend in bond Energy of the halogens above: -Decrease down the group from chlorine to Iodine Element Halogen H Na Mg Al Si C P F HF NaF MgF2 AlF3 SiF4 CF4 PF3 Cl HCl NaCl MgCl2 AlCl3 SiCl 4 CCl4 PCl3 Br HBr NaBr MgBr2 AlBr3 SiBr4 CBr4 PBr3 I Hl Nal Mgl2 All3 SiI4 C l 4 PBr3
23 -Atomic radius increase down the group decreasing the energy required to break the covalent bonds between the larger atom with reduced effective nuclear @ charge an outer energy level that take part in bonding. (c)Group VIII elements: Noble gases Group VIII elements are called Noble gases. They are all non metals. Noble gases occupy about 1.0% of the atmosphere as colourless gaseous mixture. Argon is the most abundant with 0.9%. They exists as monatomic molecules with very weak van-der-waals /intermolecular forces holding the molecules. They include: All noble gas atoms have a stable duplet(two electrons in the 1st energy level) or octet(eight electrons in other outer energy level)in the outer energy level. They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon.
They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon. The more the number of energy levels the bigger/larger the atomic size/radius. e.g. The atomic size/radius of Argon is bigger/larger than that of Neon because Argon has more/3 energy levels than Neon (2 energy levels). Atomic radius noble gases increase down the group as the number of energy levels increases. The effective nuclear attraction on the outer electrons thus decrease down the group. The noble gases are generally unreactive because the outer energy level has the stable octet/duplet. The stable octet/duplet in noble gas atoms lead to a comparatively very high 1st ionization energy. This is because losing /donating an electron from the stable atom require a lot of energy to lose/donate and make it unstable. Element Symbol Atomic number Electron structure State at room temperature Helium He 2 2: Colourless gas Neon Ne 10 2:8 Colourless gas Argon Ar 18 2:8:8 Colourless gas Krypton Kr 36 2:8:18:8 Colourless gas Xenon Xe 54 2:8:18:18:8 Colourless gas Radon Rn 86 2:8:18:32:18:8 Radioctive
24 As atomic radius increase down the group and the 1st ionization energy decrease, very electronegative elements like Oxygen and Fluorine are able to react and bond with lower members of the noble gases.e.g Xenon reacts with Fluorine to form a covalent compound XeF6.This is because the outer electrons/energy level if Xenon is far from the nucleus and thus experience less effective nuclear attraction. Noble gases have low melting and boiling points. This is because they exist as monatomic molecules joined by very weak intermolecular/van-der-waals forces that require very little energy to weaken and form liquid and break to form a gas. The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity.
The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity. Element Formula of molecule Electrical conductivity Solubility in water Atomic radius(nM) 1st ionization energy Melting point(0C) Boiling point(0C) Helium He Poor Insoluble 0.128 2372 -270 -269 Neon Ne Poor Insoluble 0.160 2080 -249 -246 Argon Ar Poor Insoluble 0.192 1520 -189 -186 Krypton Kr Poor Insoluble 0.197 1350 -157 -152 Xenon Xe Poor Insoluble 0.217 1170 -112 -108 Radon Rn Poor Insoluble 0.221 1134 -104 -93 Uses of noble gases Argon is used in light bulbs to provide an inert environment to prevent oxidation of the bulb filament Argon is used in arch welding as an insulator. Neon is used in street and advertisement light Helium is mixed with Oxygen during deep sea diving and mountaineering. Helium is used in weather balloon for meteorological research instead of Hydrogen because it is unreactive/inert. Hydrogen when impure can ignite with an explosion. Helium is used in making thermometers for measuring very low temperatures. STRUCTURE AND BONDING IONIC (ELECTROVALENT) BONDING Noble gases like neon or argon have eight electrons in their outer shells (or two in the case of helium). These noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. When other atoms react, they try to organise electrons such that their outer shells are either completely full or completely empty. Chemical reactions occur so that atoms attain inert gas configuration by either losing valency electrons as in the case of metals, or gaining electrons as in the case of non metals. Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable. Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8).
Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable. Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an electron from somewhere it too would become more stable. If a sodium atom gives an electron to a chlorine atom, both become more stable. The sodium has lost an electron, so it no longer has equal numbers of electrons and protons. Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an atom, positive ions are formed. Positive ions are sometimes called cations because they move to the cathode during electrolysis. The chlorine has gained an electron, so it now has one more electron than proton. It therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed. A negative ion is sometimes called an anion since it drifts to the anode during electrolysis. The nature of ionic bond The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges. You need one sodium atom to provide the extra electron for one chlorine atom, so they combine together 1:1. The formula is therefore NaCl. Properties of ionic compounds  All compounds with ionic bonding produce giant ionic structures.  Consist of oppositely charged ions arranged in an ionic lattice, the ions are held together by strong ionic bonds. e.g. NaCl is composed of Na+ ions and Cl- ions.  These bonds are hard to break, therefore ionic substances have very high melting and boiling points.  All exist as solids.  They conduct electricity when molten, because the ions are free to move, but do not conduct when solid.  They conduct electricity in the aqueous state because the ions are free to move.  Most ionic substances are soluble in water because the polar water molecules can accommodate the charged ions. COVALENT BONDING - SINGLE BONDS As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds.
 They conduct electricity in the aqueous state because the ions are free to move.  Most ionic substances are soluble in water because the polar water molecules can accommodate the charged ions. COVALENT BONDING - SINGLE BONDS As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds. Depending on the number of electron pairs shared between atoms which participate in bonding, covalent bonds are classified as follows: Some simple covalent molecules Chlorine For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram. The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. In reality there is no difference between them. The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms. Hydrogen Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. This is another single bond. Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure. Water Oxygen atom has six electrons in the outer shell, while each of the two hydrogen atoms has one each. After bonding, oxygen has 8 electrons while each hydrogen atom has two as shown by the molecule. NITROGEN GAS Each nitrogen atom has five electrons in the outer shell. Each needs 3 electrons to complete the outer shell. In the formation of the molecule, each nitrogen atom contributes three electrons and a triple bond is formed Characteristics of Covalent Compounds 1) Covalent compounds consist of molecules and not ions. The molecules do not have any electric charge on them. The molecules are held together by weak forces called Van der Waal's forces. 2) Covalent compounds are gases, volatile liquids or soft solids. As there are weak, Van der Waal's forces between the molecules, they are not held in rigid position. The state depends on the bond energy. If the bond energy is very low, they stay as gases, if it is appreciable they are volatile liquids.
As there are weak, Van der Waal's forces between the molecules, they are not held in rigid position. The state depends on the bond energy. If the bond energy is very low, they stay as gases, if it is appreciable they are volatile liquids. If very high, they exist as soft solids. 3) Covalent compounds generally have low melting and boiling points. As Van der Waal's forces are weak, a very small amount of energy is required to break the bond between the molecules corresponding to low melting point and boiling point. 4) Covalent compounds dissolve in organic solvents. As they do not contain ions, solvation does not take place when water is added to the compound. Hence they do not dissolve in water. 5) Covalent compounds are bad conductors of electricity. They do not contain ions in the fused state, nor do ions migrate on application of an electric potential. Hence, there is no conduction of current. 6) Covalent compounds are less dense when compared to water. Very weak Van der Waal's forces hold the molecules together, hence there are large inter molecular spaces. Consequently less number of molecules per unit volume, which means mass per unit volume is also less. Hence they have a low density. Exceptions  Diamond and graphite, the allotropes of carbon have high melting point.  Hydrogen chloride in the aqueous state conducts electricity. The giant covalent structure of diamond Carbon has an electronic arrangement of 2, 4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds. In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that's not really the case. We are only showing a small bit of the whole structure. This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal. The physical properties of diamond Diamond  Has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.  Is very hard. This is again due to the need to break very strong covalent bonds operating in 3dimensions.  Doesnಬt conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move.