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Chapter 4 – Inorganic chemistry and the periodic table Essay

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Ionic equations
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describe chemical equations by showing only the reacting ions in solutions while leaving out the spectator ions
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Spectator ions
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Ions which are present but not involved in the reaction
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Ionic precipitation reaction
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a reaction in which produces a solid precipitate on mixing 2 solutions containing ions
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Thermal decomposition
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When a compound decomposes on heating
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Thermal decomposition example and details
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green copper(II) carbonate breaks up into black copper(II) oxide and CO2. Most carbonates other than the group 1 metals decompose on heating. Hydrated compounds also decompose on heating
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Acid + metal > Acid + metal oxide > Acid + metal hydroxide > Acid + metal carbonate >
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salt + hydrogen salt + water salt + water salt + water + carbon dioxide
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Which acids are soluble/insoluble in water
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All are soluble
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Which metal hydroxides/carbonates are soluble/insoluble?
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Soluble: Alkalis, Na/K CO3 and OH. CaOH is slightly soluble.
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Which Salts are soluble/insoluble?
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Soluble: all nitrates, chlorides (except silver and lead ones), sulphates (Ca and Ag slightly)(except barium and lead aren’t)
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(Group 1) uses of lithium
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Li2CO3 is used in drugs for treating mental illness, some of its compounds are vital for organic synthesis
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(Group 1) uses of Sodium
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Sodium is a powerful reducing agent used to extract titanium, sodium vapour is used in streetlights and NaOH is the most important industrial alkali
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(Group 1) uses of Potassium
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Potassium ions are an essential nutrient in plants and an ingredient of some fertilisers
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Trend
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describes the way in which a property increases or decreases along a series of elements or compounds
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How does atomic radius change down the group
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Charge of the nucleus increases, and number of filled inner shells increases, thus the atomic radius increases down the group
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(Group 1) Why does reactivity increase down the group?
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The no. of electrons in the inner shells is always 1 less than the number of protons in the nucleus, the shielding affect increases down the group meaning that, with a constant effective nuclear charge, and the outer electrons getting further away, the outer electron is less attracted down the group.
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Alkali metal + water >
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metal hydroxide + hydrogen
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Group 1 plus chlorine…
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All G1 metals react vigorously with chlorine to for colourless ionic salts (2M + Cl2 > 2MCl) which are soluble in water
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Bases
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anti-acids, acids give up H+ ions, whilst bases accept H+ ions
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Hydroxides of G1
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all white solids, soluble in water forming alkaline solutions, solubility increases down the group. These are strong bases because they are fully ionised in water
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Carbonates of G1
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white with general formula M2CO3. They are soluble unusually and solutions of these are alkaline because the carbonate ions remove H+ ions. CO32-(aq) + H2O(l) > HCO3-(aq)+ OH-(aq). Most of these don’t decompose when heated, except for Li2CO3
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Nitrates of G1
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White crystalline solids (MNO3(s)) and highly soluble. These are much harder to decompose than most nitrates, first melting before decomposing, releasing O2 and becoming nitrites: 2KNO3(s) > 2KNO2(s) + O2(g). LiNO3 however behaves like other nitrates, decomposing to form an oxide, NO2 and O
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Why are Na and K compounds are used as chemical reagents
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the ions are unreactive and are thus spectators. Additionally their compounds are soluble in water and they are colourless so they don’t hide or interfere with colour changes
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Flame colours of G1
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lithium: bright red, sodium: bright yellow, potassium: lilac
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Why do flame colours happen?
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Flame excites outer electrons of sodium ions, raising them to higher energy levels. The atoms emit the light as the electrons drop back to lower energy levels
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Group 2 are
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Alkaline earth metals
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Difference between G1 and G2
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Harder and denser than group 1, with higher melting temperatures
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Uses of beryllium, magnesium and barium
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Be is a strong metal with a high MT, but less dense than transition metals. Mg is low density, making light alloys used in aircraft and car manufacture. Ba is soft, and highly reactive with air and moisure so is kept under oil.
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(G2) why does reactivity increase down the group?
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The increasing no. of shells down the group means the atomic radii increase down the group hence the reactivity increases down the group as the pull on outer electrons decreases
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G2 ionisation energy trends
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The 1st and 2nd decrease down the group, the shielding means the greater effective nuclear charge attracting the outer electron is 2+. Down the group the outer electrons get further away from the same nuclear charge so they are attracted less
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(Group 2) Mg reacting with oxygen
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Mg burns with a white flame, which is why it is used in fireworks and flares. It produces MgO
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(Group 2) Ca reacting with oxygen
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Ca burns with a red flame forming CaO, strontium reacts similarly
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(Group 2) Ba reacting with oxygen
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Ba burns with a green flame to form Barium peroxide, BaO2
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(Group 2) Mg reacting with water
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Mg reacts very slowly with cold water producing Mg(OH)2 and H2. It reacts better with steam
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(Group 2) Ca reacting with water
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Ca reacts to produce H2 and Ca(OH)2, initially the Ca(OH)2 dissolves but the solubility is low so as more forms the solution becomes saturated and a precipitate appears
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(Group 2) Ba reacting with water
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Ba reacts even faster than Ca and its hydroxide is more soluble
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How does group 2 react with chlorine?
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all G2 metals react with Cl on heating to form white chlorides, MCl2
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Basic oxide
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a metal oxide which reacts to form salts and water. Soluble basic oxides are alkalis as the oxide acts as a base taking the H+ ions from the acid. All G2 metals apart from Be form these.
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Magnesium oxide
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MgO is a white solid, which turns into a hydroxide when added to water. This is slightly soluble. MgO has a high melting point so is often used as a heat resistant ceramic to line furnaces
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Calcium oxide
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CaO is a white solid made by heating CaCO3, it reacts vigorously with cold water, hence the traditional name ‘quicklime’, and this reaction produces calcium hydroxide.
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Magnesium hydroxide
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Mg(OH)2 is the active ingredient in milk of magnesia, used as an antacid and a laxative
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Calcium hydroxide
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Ca(OH)2 is slightly soluble in water forming an alkaline solution, limewater. The CO2 test works because a solution of Ca(OH)2 reacts with the gas forming a white insoluble calcium carbonate precipitate
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Barium hydroxide
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Ba(OH)2 is soluble, sometimes used as an alkali in chemical analysis. It is better than sodium and potassium for this as it can’t be contaminated by its carbonate, because barium carbonate is insoluble in water
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G2 hydroxides
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All have the formula M (OH)2 and are to varying degrees soluble in water, forming alkaline solutions, however their solubility increases down the group
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G2 carbonates
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They all have the formula MCO3 and are insoluble in water, they all react with dilute acids and decompose on heating to give the oxide and CO2. However they are more difficult to decompose down the group – they become more thermally stable
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G2 nitrates
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They all have the formula M(NO3)2, are colourless, crystalline solids, very soluble in water and decompose to the oxide on heating. They become more difficult to decompose down the group
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G2 sulphates
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All colourless solid with the formula MSO4 but are less soluble down the group
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Flame colours for G2
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beryllium and magnesium – no colour, calcium – brick red, strontium – bright red, barium – pale green
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Thermal stability
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an indication as to the ease with which compounds decompose on heating. Compounds are stable if they do not tend to decompose into their elements or into other compounds. A compound, which is stable at RTP may become more or less stable as conditions change
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Polarising power
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an indication of the extent to which a + ion is able to distort the electron cloud around a – ion. The larger the charge of the + ion and the smaller its size, the greater its polarising power
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G2 carbonates and nitrates are generally less stable than the corresponding G1 compounds. This suggests…
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…that the larger the charge on the metal ion, and the smaller the metal ion, the less stable the compounds
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Trend in stability in carbonates..?
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The carbonates become more stable down group 1 and 2, this helps confirm that the larger the metal ion, the more stable the compound
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thermal stability explained in terms of polarising power
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The larger the charge and the smaller the ionic radius, the greater the charge density and thus the greater the polarising power of the ion. Higher polarising power means the ions pulls at the electrons more, distorting the bonding and making it easier to break up the – ion into an oxide ion and CO2
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why does thermal stability increase down the group?
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Although the charge remains constant down the group the size increases, decreasing the charge density, and thus the polarising power. Thus the thermal stability of the nitrates and carbonates increases down the group
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Uses of the halogens
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They are used as ingredients in plastics, pharmaceuticals, anaesthetics and dyestuffs
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Appearance of halogens
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At RTP, chlorine is a yellow-green gas, bromine is a dark red liquid and iodine is a dark grey solid. They exist as diatomic molecules linked by a single covalent bond.
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Halogens and intermolecular forces
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Intermolecular forces increase down the group as the numbers of electrons in the molecules increase, the larger molecules are therefore more polarisable so M and B temps increase down the group.
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Fluorine
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most electronegative element, oxidation state -1
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Chlorine
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1 oxidation state, but can be oxidised to positive states by F and O. Chlorine is used in the product of polymers (PVC). It is also used to ‘chlorinate’ drinking water.
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Bromine
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Reactive but a less powerful oxidising agent than fluorine, it is used to make medicines, dyes and flame-retardants.
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Iodine
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Used to make medicines, dyes and catalysts. Iodine is important in the diet so that the thyroid gland can make the hormone thyroxine, which regulates growth and metabolism.
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Halogens and hydrocarbon solvents
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Dissolve in hydrocarbon solvents (cyclohexane), when dissolved, the solution will appear a similar colour to the free halogen vapours (iodine – violet…)
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Halogens and water
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Halogens are less soluble in water, Cl and Br(aq) are useful reagents, colours are similar to those of vapours. These also react with water. I doesn’t dissolve in water but is does dissolve in KI(aq), I2 reacts with the iodide ions to form triiodide ions (yellow-brown colour)
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(G7) Cl and Br + metals
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react with G1 and 2 to form ionic halides
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(G7) I + metals
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reacts to form iodides but because of the polarisability of the iodide ion, iodides formed with small cations (Li+) or highly charged ones (Al3+) are essentially covalent
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(G7) Halogens + d block metals
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Halogens do react with most d block metals – when they react with hot iron, they form Iron (III) chloride, Iron(III) bromide and iron(II) iodide as iodine reduces the iron (III) ions
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(G7) Chlorine + non metals
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Cl reacts with most non metals to form molecular chlorides, e.g. SiCl4(l). However chlorine will not react directly with C, O or F. H burns in Cl to produce colourless HCl(g)
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(G7) Bromine + non metals
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Bromine oxidises non-metals such as S and H when heated forming molecular bromides. Bromine vapour and hydrogen gas react smoothly with a pale bluish flame
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(G7) Iodine + non metals
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Iodine oxidises H on heating to form HI, which is a reversible reaction
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Halogens + Fe2+ ions (aq)
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Cl and Br can oxidise iron (II) ions(aq) but iodine cannot
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AgCl: 1) obs when AgNO3 is added, 2) obs when ammonia is added
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1) White precipitate turns purple-grey in sunlight 2)Dissolves in dilute ammonia, forming a colourless solution
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AgBr: 1) obs when AgNO3 is added, 2) obs when ammonia is added
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1) Creamy precipitate 2)Dissolves in concentrated ammonia forming a colourless solution.
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AgI: 1) obs when AgNO3 is added, 2) obs when ammonia is added
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1)Yellow precipitate 2)Does not dissolve.
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AgF: 1) obs when AgNO3 is added, 2) obs when ammonia is added
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Silver fluoride is soluble so there is no precipitate to start with.
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Halogens reacting with concentrated sulphuric acid overview..
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NaCl + NaSO4 -> HCl (g). The gas fumes in moist air. This doesn’t work with Br as bromide ions are oxidised to bromine by sulphuric acid. Iodide ions as such strong reducing agents that they reduce sulphur from +6 to 0 and -2
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NaCl reacting with concentrated H2SO4 (Halide Cl-).
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> HCl + NaHSO4. Product: HCl, observation: gas which fumes in moist air, reaction type: acid -base.
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Bromide reacting with H2SO4. Product: Br2
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Observation: orange vapour, reaction type: oxidation, equation: 2Br- > Br2 + 2e-
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Bromide reacting with H2SO4. Product: SO2
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Observation: Colourless acidic gas, Reaction type: Reduction, Equation: H2SO4 + 2H+ + 2e- > SO2 + 2H2O
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Bromide reacting with H2SO4. Product: HBr
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Observation: Steamy fumes, Reaction type: Acid-base, Equation: 2NaBr + H2SO4 > HBr + NaHSO4
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Iodide reacting with H2SO4. Product: HI
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Observation: Steamy fumes, Reaction type: Acid-base Equation: NaI + H2SO4 > HI + NaHSO4
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Iodide reacting with H2SO4. Product: I2
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Observation: Purple solid, Reaction type: Oxidation, Equation: H2SO4 + 8H+ + 8I- > 4I2 + H2S + 4H2O
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Iodide reacting with H2SO4. Product: H2S
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Observation: – Reaction type: Reduction, Equation: H2SO4 + 8H+ + 8I- > 4I2 + H2S + 4H2O
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Iodide reacting with H2SO4. Product: SO2
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Observation: Gas with bad egg smell, Reaction type: Reduction, Equation: H2SO4 + 2H+ + 2I- > I2 + SO2 + H2O
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Iodide reacting with H2SO4. Product: S
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Observation: Yellow solid, Reaction type: Reduction, Equation: H2SO4 + 6H+ + 6I- > 3I2 + S + 4H2O
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Halide reducing and oxidising trends
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The power of halides as reducing agents increases down the group, and for oxidising agents increases up the group.
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Hydrogen halides
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They are colourless polar, molecular compounds at RT that fume in moist air, with formula HX. They are soluble in water, forming acidic solutions, which ionise completely in water.
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hydrogen halide + ammonia
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Mixing the hydrogen halides with ammonia produces a white smoke of an ammonium salt. Ammonia molecules become ammonium ions in this reaction
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Oxoanion
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an ion with the general formula XxOyz- (where X is any element and O is oxygen)
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What happens when halogens react with water and alkalis
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Oxoanions form. Cl dissolves in water to form a mix of weak chloric and strong hydrochloric acid. This is a disprop reaction. Cl2 + H2O > HCl + HClO Br reacts similarly but to a lesser extent and I is insoluble in water and barely reacts
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(G7 in +1 and +5 states) When chlorine dissolves in KOH at RT, it produces
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chlorate (I) and chloride ions: The active ingredient in bleach is sodium chlorate, made by using NaOH in this reaction. When heated the chlorate (I) ions disproportionate to chlorate (V) and chloride ions. Cl2 + 2OH- > ClO + Cl + H2O
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(G7 in +1 and +5 states) When Br and I react with alkalis
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Br and I react in a similar way to Cl with alkalis, though the BrO- and IO- are less stable, so they disproportionate at a lower temperature.