# WJEC Chemistry: 1.6- The Periodic Table

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• Created on: 10-05-16 18:34

## Periodicity: Across a Period

Trends:

• Atomic Size (DECREASES)-  The nuclear charge increases but the shielding stays the same, therefore the electrons are attracted closer to the nucleus.
• Ionisation Energies (GENERALLY INCREASES)- The nuclear charge increases but the shielding remains the same, making the electrons harder to remove.                                                                       - The I.E decreases from Group II to Group III. This is because the valence electron is in a new subshell.                                                                                                                                                                 - The I.E decreases from Group V to Group VI because the valence electron becomes paired and the repulsion in the orbital makes it easier to remove.
• Electronegativity (INCREASES)- The nuclear charge increases but the shielding remains the same, making the electrons more strongly attracted to the atom, so that atoms will have a larger share of the electrons in a covalent bond.
• Melting Temperature (DEPENDS ON BONDS)- The structure and bonding of the elements in a period varies widely. The trends in the intramolecular bonding causes a difference in physical properties i.e their melting points.
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## Periodicity: Down a Group

Trends:

• Ionisation Energies (DECREASES)- The valence electrons are further away from the nucleus making it easier to remove, and the nucleus is more effectively shielded by more inner shells of electrons.
• Electronegativity (DECREASES)- As the atoms get bigger there is more shielding, resulting in a reduced ability to attract the electron density in a covalent bond towards itself.
• Atomic Size (INCREASES)- Each element has an extra filled main level of electrons, resulting in the valence electrons becoming further away from the nucleus.
• Melting Temperature (DEPENDS ON BONDS)- Depends on the group and the type of bonding in the element found within the group.                                                                                                               - Metallic Elements (Groups 1/2): DECREASES because the smaller atoms have a larger charge density compared to larger atoms, resulting in the strength of metallic bonds decreasing and less energy required to break these bonds.                                                                                                           - Non-Metallic Elements (Groups 5-8): INCREASES because larger atoms have more electrons making their Van der Waals between the molecules stronger. Resulting in more energy required to break these bonds.
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## Group 2: Reactions with Oxygen and Water

Reactions with Oxygen:

• All the metals react with oxygen to form a solid ionic metal oxide. They become more reactive going down the group because as it becomes easier for the atoms to lose their 2 valence electrons to form a +2 ion.                                                                                                                               - e.g. 2Mg(s) + O2(g) --> 2MgO(s)

Reactions with Water

• Metals react with water to form metal hydroxides and hydrogen gas. The reactivity to water increases down the group as electrons are more easily removed. Beryllium and Magnesium do not react to water. However, all other Group 2 metals react with water.                                                        - Ca(s) + 2H2O(l) --> Ca(OH)2(s) + H2(g)
• Metals react with steam to form metal oxides and hydrogen gas. The reactivity to steam increases down the group as electrons are easier to remove. Beryllium does not react with water or steam, Magnesium does not react with water but reacts with steam. All other Group 2 metals react with steam.                                                                                                                                                                   - Mg(s) + H20(g) --> MgO(s) + H2(g)
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## Group 2: Reactions with Ions

Reactions with Hydroxides:

• Magnesium Hydroxide is insoluble, so will form a white precipitate with a hydroxide solution. Calcium Hydroxide is sparingly soluble, so will form a faint white precipitate with a hydroxide solution. Barium Hydroxide is soluble so no reaction happens between the solutions.                         - Ca2+(aq) + 2OH-(aq) --> Ca(OH)2(s)

Reactions with Carbonate:

• All the metal ions (Mg2+, Ca2+ and Ba2+) will react with an aqueous solution of a carbonate to form a white precipitate, as they are all insoluble.                                                                                       - Mg2+(aq) + CO32-(aq) --> MgCO3(s)

Reactions with Sulphate:

• Magnesium Sulphate is soluble so no reaction will happen between the two solutions. Calcium Sulphate is sparingly soluble, so a faint white precipitate of calcium sulphate will be observed on the addition of a sulphate solution. Barium Sulphate is insoluble, so will form a thick white precipitate on the addition of a soluble solution of sulphate ions.                                                             - Ba2+(aq) + SO42-(aq) --> BaSO4(s)
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## Group 2: Solubility and Uses

Solubility of Sulphates:

• The solubility of the sulphates decreases down Group II. MgSO4 is soluble, CaSO4 is sparingly soluble and SrSO4 and BaSO4 are insoluble. If sulphuric acid or sodium sulphate is added to the aqueous solutions of calcium, strontium or barium ions, a white precipitate will be formed.              - Ba2+(aq) + SO42-(aq) --> BaSO4(s)

Solubility of Hydroxides:

• The solubility of the Group II Hydroxides increases down the group. Mg(OH)2 is insoluble, Ca(OH)2 is sparingly soluble and Sr(OH)2 and Ba(OH)2 are soluble.                                                        - Mg2+(aq) + 2OH-(aq) --> Mg(OH)2(s)

Uses of Sulphates and Hydroxides:

• Magneisum Hydroxide is almost insoluble and is sold as a suspension in water. It is known as "milk of magnesia" and is taken to alleviate constepation.
• Barium Sulphate can be eaten as part of a "barium meal". This is because Barium is good at absorbing X-Rays and when the Barium Sulphate gets to the gut the outline can be located using X-Rays.
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## Group 2: Thermal Stability

Both the Carbonates and Hydroxides thermal stability increases down the group because of the increase in ionic radius of the cation and its polarizing power decreasing.

Thermal Stability of Carbonates:

• Group II carbonates thermally decompose on heating to produce their corresponding oxide. The carbonates become increasingly stable down the group, so more heat energy is needed to break them down.                                                                                                                                                         - MgCO3(s) --> MgO(s) + CO2(g)

Thermal Stability of Hydroxides:

• Group II hydroxides thermally decompose on heating to produce their corresponding oxide. The hydroxides become increasingly stable down the group, so more heat energy is needed to break them down.                                                                                                                                                         - Ca(OH)2(s) --> CaO(s) +H2O(g)

General Trend in Reactivity (Group I & II):

• Reactivity increases down the the groups. This is because the valence electrons are further away from the nucleus so they are easier to lose, meaning that elements with bigger atomic radii more reactive.
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## Group 7: Trends in Reactivity

The general trend for the reactivity of the halogens is that it decreases down the group. This is because as the atoms get bigger the valence electrons are attracted less strongly to the nucleus due to more efficient shielding.

Volatility:

• The volatility decreases down the group. This is because the number of electrons increases, therefore the number of Van der Waals forces increases and resulting in the volatility decreasing.

Oxidising Ability:

• The halgoen molecules are good oxidising agents, readily gaining electrons to form the corresponding halide ion. Flourine is the most reactive halogen and the strongest oxidising agent in the group.                                                                                                                                                         - F2 + 2e- --> 2F-
• The oxidising ability decreases down the group. This is because as the atoms get bigger, valence electrons are attracted less strongly to the positive nucleus.
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## Group 7: Metal and Displacement Reactions

Reactions with Metals:

• The halogens react with metals to form salts. In both these reactions the metal is oxidised and the halogen is reduced. These reactions are known as redox reactions.                                                         - Iron + Chlorine --> Iron(III)Chloride   or   2Fe(s) + 3Cl2(g) --> 2FeCl3(s)                                               - Sodium + Iodine --> Sodium Iodide   or   2Na(s) + I2(g) --> 2NaI(s)

Displacement Reactions:

• Adding solutions containing the halogen elements to solutions of other halide ions can help determine which halogen is the most reactive and the strongest oxidising agent.                                 - Cl₂(aq) + 2NaBr(aq) --> Br₂(aq) +2NaCl(aq)   or   Cl₂(aq) + 2Br⁻(aq) --> Br₂(aq) + 2Cl⁻(aq)
• All halogen displacement reactions are redox reactions and the more reactive halogen will oxidise the halide ion of a less reactive halogen. Because Fluorine is the most reactive halogen, it will displace all halogens bar itself to their original solutions. Therefore, because Iodine is the least reactive it will not displace any of the halogens.
• The colour of the halogen solutions that the displacement reactions revert to is: Fluorine (Pale Yellow), Chlorine (Yellow-Green), Bromine (Red-Brown) and Iodine (Dark Brown)
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## Group 7: Testing for Halide ions and Uses

Testing for Halide Ions:

• Aqueous solutions of chloride, bromide and iodide ion can be tested for and identified using silver ions that that have been dissolved in dilute nitric acid. Dilute nitric acid must first be added to remove any other ions i.e hydroxides and carbonates present as these anions also produce precipitates with silver ions and may intefere with the test.                                                                      - AgNO3(aq) + Cl-(aq) --> AgCl(s) + NO3-(aq)   or   Ag+(aq) + Cl-(aq) --> AgCl(s)
• The precipitates formed are: Flouride (None), Chloride (White), Bromide (Cream) and Iodide (Yellow).
• Ammonia solution is then added to distinguish the presence of the anion. Silver Chloride dissolves in dilute ammonia solution but Silver Bromide and Silver Iodide are both insoluble so their precipitates remain.

Uses of Chlorine and Fluoride in Water Treatment:

• Chlorine is used in water treatment and makes water safe to drink and use by killing pathogenic bacteria and viruses and preventing the outbreak of serious diseases such as typhoid and choleraFluoride is generally added to water and toothpaste to reduce tooth decay caused by cavity formation and in strengthening bones, thus reducing osteoperosis although too much fluoride can harm tooth development.
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