inorganic chem

trends in atomic radius , from Na-Ar

Atomic radii decrease from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.

trends in 1st ionisation energy , from Na-Ar

There is a general trend across to increase. This is due to increasing number of protons as the electrons are being added to the same shell. There is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al's electron is slightly easier to remove because the 3p electrons are higher in energy. There is a small drop between phosphorous and sulfur. Sulfur's outer electron is being paired up with an another electron in the same 3p orbital. When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.

trends in melting points , from Na-Ar

For Na, Mg, Al- Metallic bonding : strong bonding - gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller sized ion with a greater positive charge also makes the bonding stronger. Higher energy is needed to break bonds. Si is Macromolecular: many strong covalent bonds between atoms, high energy needed to break covalent bonds- very high mp +bp Cl2 (g), S8 (s), P4 (S)- simple molecular : weak van der waals between molecules, so little energy is needed to break them - low mp+ bp S8 has a higher mp than P4 because it has more electrons (S8 =128)(P4=60) so has stronger v der w between molecules Ar is monoatomic weak van der waals between atoms

trends in atomic radius , from Mg- Ba

Atomic radius increases down the group. As one goes down the group, the atoms have more shells of electrons making the atom bigger

trends in 1st ionisation energy, from Mg- Ba

The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells. In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons.

trends in melting points , from Mg- Ba

Melting points decrease down the group. The metallic bonding weakens as the atomic size increases. The distance between the positive ions and delocalized electrons increases. Therefore the electrostatic attractive forces between the positive ions and the delocalized electrons weaken.

magnesiums reaction with steam and water

Magnesium reacts in steam to produce magnesium oxide and hydrogen. The Mg would burn with a bright white flame. The MgO appears as a white powder. Mg(s) + H2O(g) MgO(s) + H2 (g)

Mg will also react with warm water, giving a different magnesium hydroxide product. Mg + 2 H2O Mg(OH)2 + H2 This is a much slower reaction than the reaction with steam and there is no flame

trends in the reaction with water , from Mg- Ba

The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides. Ca + 2 H2O (l) Ca(OH)2 (aq) + H2 (g) Sr + 2 H2O (l) Sr(OH)2 (aq) + H2 (g) Ba + 2 H2O (l) Ba(OH)2 (aq) + H2 (g)

One would observe: fizzing, (more vigorous down group) the metal dissolving, (faster down group) the solution heating up (more down group) with calcium a white precipitate appearing (less precipitate forms down group with other metals)

use of magnesium in titanium extraction from TiCl4

eps in extracting titanium 1. TiO2 (solid) is converted to TiCl4 (liquid) at 900C: 2. The TiCl4 is purified by fractional distillation in an argon atmosphere. 3. The Ti is extracted by Mg in an argon atmosphere at 500C

TiO2 + 2 Cl2 + 2 C TiCl4 + 2 CO

TiCl4 + 2Mg Ti + 2 MgCl2

trends in relative solubilities of hydroxides in water , from Mg- Ba

Group II hydroxides become more soluble down the group. All Group II hydroxides when not soluble appear as white precipitates

Magnesium hydroxide is classed as insoluble in water. Simplest ionic equation for formation of Mg(OH)2 (s) Mg2+(aq) + 2OH-(aq) Mg(OH)2 (s).

Calcium hydroxide is classed as partially soluble in water and will appear as a white precipitate

Barium hydroxide would easily dissolve in water. The hydroxide ions present would make the solution strongly alkaline

Use of Mg(OH)2 in medicine?

Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.

The use of Ca(OH)2 in agriculture

e It is used in agriculture to neutralise acidic soils.

The use of CaO or CaCO3 to remove SO2 from flue gases

Calcium oxide can be used to remove SO2 from the waste gases from furnaces (e.g. coal fired power stations) by flue gas desulfurization. The gases pass through a scrubber containing basic calcium oxide which reacts with the acidic sulfur dioxide in a neutralisation reaction. The calcium sulfite which is formed can be used to make calcium sulfate for plasterboard. SO2 + CaO CaSO3 calcium sulfite

trends in solubilities of sulfates in water, from Mg- Ba

Group II sulfates become less soluble down the group.

use of BaCl2 solution to test for sulfate ions

BaCl2 solution acidified with hydrochloric acid is used as a reagent to test for sulfate ions. If acidified barium chloride is added to a solution that contains sulfate ions a white precipitate of barium sulfate forms.

The use of BaSO4 in medicine

BaSO4 is used in medicine as a 'Barium meal' given to patients who need x-rays of their intestines. The bariumabsorbs the x-rays and so the gut shows up on the x-ray image. Even though barium compounds are toxic, it is safe to use here because barium sulfate low solubility means it is not absorbed into the blood.

the trends in electronegativity of the halogens

Electronegativity decreases down the group as there are more shells leading to more shielding and electrons being held at a further distance.

trends in boiling points of the halogens

Increase down the group As the molecules become larger they have more electrons and so have larger van der waals forces between the molecules. As the intermolecular forces get larger more energy has to be put into break the forces. This increases the melting and boiling points.

trend in oxidising ability of halogens down the group

A halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds. The oxidising strength decreases down the group. Oxidising agents are electron acceptors.

reactions of solid sodium F- and cl ions with concentrated sulfuric acid

NaF(s) + H2SO4 (l) NaHSO4 (s) + HF(g) Observations: White steamy fumes of HF are evolved. NaCl(s) + H2SO4 (l) NaHSO4 (s) + HCl(g) Observations: White steamy fumes of HCl are evolved.

F- and Clions are not strong enough reducing agents to reduce the S in H2SO4 . No redox reactions occur. Only acid-base reactions occur.

reactions of solid sodium bromide ions with concentrated sulfuric acid

Br- ions are stronger reducing agents than Cl- and F- and after the initial acid-base reaction, the bromide ions reduce the sulfur in H2SO4 from +6 to + 4 in SO2 Bromide Acid- base step: NaBr(s) + H2SO4 (l) NaHSO4 (s) + HBr(g) Redox step: 2 H+ + 2 Br - + H2SO4 Br2 (g) + SO2 (g) + 2 H2O(l)

Ox ½ equation 2Br - Br2 + 2e- Re ½ equation H2SO4 + 2 H+ + 2 e- SO2 + 2 H2O

Observations: White steamy fumes of HBr are evolved. orange fumes of bromine are also evolved and a colourless, acidic gas SO2

reactions of solid sodium iodide ions with concentrated sulfuric acid

I- ions are the strongest halide reducing agents. They can reduce the sulfur from +6 in H2SO4 to + 4 in SO2 , to 0 in S and -2 in H2S. NaI(s) + H2SO4 (l) NaHSO4 (s) + HI(g) 2 H+ + 2 I- + H2SO4 I2 (s) + SO2 (g) + 2 H2O(l) 6 H+ + 6 I- + H2SO4 3 I2 + S (s) + 4 H2O (l) 8 H+ + 8 I- + H2SO4 4 I2 (s) + H2S(g) + 4 H2O(l) Observations: White steamy fumes of HI are evolved. Black solid and purple fumes of Iodine are also evolved A colourless, acidic gas SO2 A yellow solid of sulfur H2S (Hydrogen sulfide), a gas with a bad egg smell, Ox ½ equation 2I - I2 + 2e- Re ½ equation H2SO4 + 2 H+ + 2 e- SO2 + 2 H2O Re ½ equation H2SO4 + 6 H+ + 6 e- S + 4 H2O Re ½ equation H2SO4 + 8 H+ + 8 e- H2S + 4 H2O Reduction product = sulfur dioxide. Note the H2SO4 plays the role of acid in the first step producing HI and then acts as an oxidising agent in the three redox steps. Reduction products = sulfur dioxide, sulfur and hydrogen sulfide

use of acidified silver nitrate solution to identify and distinguish between halide ions

made acidic with nitric acid,

Fluorides produce no precipitate Chlorides produce a white precipitate Ag+ (aq) + Cl - (aq) AgCl(s) Bromides produce a cream precipitate Ag+ (aq) + Br- (aq) AgBr(s) Iodides produce a pale yellow precipitate Ag+ (aq) + I- (aq) AgI(s)

solubility of silver halides with ammonia

Silver chloride dissolves in dilute ammonia to form a complex ion AgCl(s) + 2NH3 (aq) [Ag(NH3 )2 ]+ (aq) + Cl - (aq) Colourless solution Silver bromide dissolves in concentrated ammonia to form a complex ion AgBr(s) + 2NH3 (aq) [Ag(NH3 )2 ]+ (aq) + Br - (aq) Colourless solution Silver iodide does not react with ammonia - it is too insoluble.

reaction of chlorine with water to form chloride ions and chlorate

Cl2 (g) + H2O (l) ⇌ HClO (aq) + HCl (aq)

(disproprotionation)

reaction of chlorine with water to form chloride ions and oxygen

If the chlorine is bubbled through water in the presence of bright sunlight a different reaction occurs. 2Cl2 + 2H2O 4H+ + 4Cl - + O2 The same reaction occurs to an equilibrium mixtureof chlorine water when standing in sunlight. The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O2 ) is produced.

pros and cons of chlorinating water

Chlorine is used in water treatment to kill bacteria. It has been used to treat drinking water and the water in swimming pools. The benefits to health of water treatment by chlorine outweigh its toxic effects.

reaction of chlorine with cold dilute aqueous NaOH and uses of product

Cl2 ,(and Br2 , I2 ) in aqueous solutions will react with cold sodium hydroxide. The colour of the halogen solution will fade to colourless. Cl2 (aq) + 2 NaOH (aq) NaCl (aq) + NaClO (aq) + H2O (l) The mixture of NaCl and NaClO is used as bleach and to disinfect/ kill bacteria.

required practical 4 tests for group 2 metals

Four test tubes should be placed in a test tube rack

Around 10 drops of 0.1 mol dm-3 barium chloride solution should be added to the first test tube

Around 10 drops of dilute sodium hydroxide solution (NaOH) should be added to the same test tube

Swirl the test tube carefully to mix well

Continue to add sodium hydroxide dropwise to the test tube, until it is in excess

This should then be repeated in the other test tubes, for calcium bromide solution, magnesium chloride solution and strontium chloride solution. The same test as above can also be done using ammonia solution and sulfuric acid solution.

colourless solution produces slight white precipitate with calcium bromide, magneisum chloride and strontium chloride. nothing happens with barium chloride. the precipitate in magnesium chloride intensifies as more NaOH is added

required practical 4 test for ammonium ions

About 10 drops of a solution containing ammonium ions, such as ammonium chloride, should be added to a clean test tube

About 10 drops of sodium hydroxide should be added using a pipette

The test tube should be swirled carefully to ensure that it is mixed well

The test tube of the solution should then be placed in a beaker of water, and the beaker of water should be placed above a Bunsen burner, so that it can become a water bath

As the solution is heated gently, fumes will be produced

A pair of tongs should be used to hold a damp piece of red litmus paper near the mouth of the test tube, to test the fumes

The red litmus paper will change colour and become blue in the presence of ammonia gas

test for halide ions

The sample being tested should be added using a pipette to a test tube

The test tube should be placed into a test tube rack

A small amount of nitric acid should be added to the sample using a pipette, followed by a small amount of silver nitrate solution

A precipitate will form, either white, cream or yellow, if a halide ion is present in the sample

The white precipitate will form if chloride ions are present in the sample

The white precipitate is AgCl

The cream precipitate will form if bromide ions are present in the sample

The cream precipitate is AgBr

The yellow precipitate will form if iodide ions are present in the sample

The yellow precipitate is AgI

further ammonia tests

add more dilute then concentrated ammonia to see what dissolves and when

test for hydroxide ions

A small amount (around 1 cm3) of the solution should be added to a test tube using a pipette

Test the pH of the solution using red litmus paper or universal indicator paper

The presence of hydroxide ions will turn the red litmus paper blue and the pH will be clearly alkaline on the universal indicator paper if hydroxide ions are present

test for carbonate ions

A small amount (around 1 cm3) of dilute hydrochloric acid should be added to a test tube using a pipette

An equal amount of sodium carbonate solution should then be added to the test tube using a clean pipette

As soon as the sodium carbonate solution is added, a bung with a delivery tube should be attached to the test tube

The delivery tube should transfer the gas which is formed into a different test tube which contains a small amount of limewater (calcium hydroxide solution)

Carbonate ions will react with hydrogen ions from the acid to produce carbon dioxide gas

Carbon dioxide gas will turn the limewater milky

test for sulfate ions

Acidify the sample with dilute hydrochloric acid and then add a few drops of aqueous barium chloride

If a sulfate is present then a white precipitate of barium sulfate is formed:

Ba2+ (aq) + SO42- (aq) → BaSO4 (s)

appearance of the halogens

Fluorine (F2 ): very pale yellow gas. It is highly reactive Chlorine : (Cl2 ) greenish, reactive gas, poisonous in high concentrations Bromine (Br2 ) : red liquid, that gives off dense brown/orange poisonous fumes Iodine (I2 ) : shiny grey solid sublimes to purple gas.

displacement reactions of halide ions

chlorine displaces all potassium halides. bromine only displaces iodide, iodine causes on change