G.C.E. A-Level Chemistry Grade 12 Unit 6

The tendency of a carbonate compound to decompose into a metal oxide and carbon dioxide gas, with the decomposition temperature increasing as you move down the group.

The ability of a compound to dissolve in water, which varies with the compound and the size of the cation. Group 2 salts with uninegative anions are generally soluble, while those with multi-negative anions tend to be insoluble.

The solubility of Group 2 hydroxides increases as you move down the group, meaning they become more soluble. This is because the smaller cations have a stronger attraction to the hydroxide anions.

Group 2 sulphates become less soluble as you move down the group, meaning they become less soluble. This is because the larger cations have a weaker attraction to the sulphate anion.

Group 2 hydrogen carbonates (bicarbonates) are only stable in aqueous solutions. In solid form, they decompose at room temperature.

Lithium, sodium, and potassium react vigorously with dilute acids to produce hydrogen gas and the corresponding metal salt.

Group 1 nitrates decompose upon heating, producing different products depending on the metal. LiNO3 produces lithium oxide, nitrogen dioxide, and oxygen. Other nitrates produce metal nitrite and oxygen.

The ability of a compound to resist decomposition when heated.

Group 1 carbonates decompose upon heating to form metal oxides and carbon dioxide. Li2CO3 is less stable than other carbonates and decomposes readily.

Group 1 bicarbonates decompose upon heating, producing metal carbonates, water vapor, and carbon dioxide.

Group 1 hydroxides react with carbon dioxide to form carbonates. These carbonates can further react with carbon dioxide to form metal hydrogen carbonates.

Group 1 elements react with hydrogen to form solid, ionic metal hydrides. These hydrides react vigorously with water producing hydrogen gas.

Different forms of the same element having different physical and chemical properties, such as rhombic and monoclinic sulphur.

A crystalline form of sulphur stable below 95°C, consisting of S8 rings packed in a particular arrangement.

A crystalline form of sulphur stable above 95°C, consisting of S8 rings packed in a different arrangement than rhombic sulphur.

A non-crystalline, elastic form of sulphur formed by rapid cooling of molten sulphur.

The ability of a substance to act as both an acid and a base.

The process where water molecules donate and accept protons, resulting in the formation of hydronium (H3O+) and hydroxide (OH-) ions.

A compound that can donate a proton (H+) to a base or accept a proton from an acid.

A nonplanar molecule with two OH groups that do not lie in the same plane, resulting in a bent structure with a dihedral angle between the OH groups.

Elements in the fourth period of the d block (except Sc and Zn) can form cations with multiple oxidation states.

The number of d electrons participating in bonding determines the oxidation state.

The 4s orbital has a lower effective nuclear charge than the 3d orbital, meaning its electrons are easier to remove during ionization.

The maximum possible oxidation state for a d-block element is the sum of its 4s and 3d electrons.

The highest possible oxidation state is typically achieved by losing all 4s and 3d electrons.

As you move across the fourth period (from left to right), the nuclear charge increases, making the 3d orbitals more stable and harder to remove electrons from.

Transition metals can act as both oxidizing and reducing agents due to their ability to interconvert between oxidation states.

Copper (Cu) commonly forms +1 oxidation state due to its 3d104s1 configuration, where it easily loses the single 4s electron.

Ammonia acts as a Brønsted-Lowry base, accepting a proton from water to form ammonium ions (NH4+) and hydroxide ions (OH-).

Ammonium salts react with strong bases like NaOH to produce ammonia gas, water, and the corresponding salt. This reaction is used to detect ammonia.

Ammonia (NH3) can act as a reducing agent with chlorine (Cl2), resulting in different products depending on the excess reagent. With excess ammonia, nitrogen gas (N2) is produced. With excess chlorine, nitrogen trichloride (NCl3) is formed.

Nitrogen trichloride (NCl3) is a covalent chloride that reacts with water to produce ammonia and hypochlorous acid (HOCl). This property makes NCl3 useful in water disinfection.

Gaseous ammonia reacts with hydrogen chloride (HCl) to form a white smoke of solid ammonium chloride (NH4Cl). This reaction is a confirmation test for ammonia presence.

Ammonium salts decompose upon heating, releasing ammonia gas and the corresponding acidic gas. This is a general property of ammonium salts.

Some ammonium salts decompose into specific products upon heating. For example, NH4NO2 decomposes into N2 and H2O. The products vary depending on the anion in the salt.

Ammonia can act as an oxidizing agent and an acid under dry conditions, reacting with metals like sodium (Na) and magnesium (Mg) to form amides (NaNH2) and nitrides (Mg3N2), respectively.

Carbon monoxide (CO) is commonly used as a reducing agent in industrial processes. It acts as a crucial component in various catalytic reactions, thanks to its lone pair of electrons on the carbon atom, making it an effective ligand.

Carbon dioxide (CO2) solidifies when subjected to low temperatures or high pressures. This solid form, known as dry ice, sublimates (transforms directly from solid to gas) under normal atmospheric conditions, making it useful as a coolant in various applications.

Carbonic acid (H2CO3) is a weak acid formed by dissolving carbon dioxide (CO2) in water. Its structure features a carbon atom bonded to two oxygen atoms and two hydroxyl groups.

Nitrogen, the first element in Group 15, exhibits distinct characteristics compared to other elements within the group. Nitrogen is a nonmetal with oxidation states ranging from -3 to +5, while other Group 15 elements exhibit metallic properties.

The metallic character of elements in Group 15 increases as you move down the group. This trend arises from increasing atomic size and decreasing electronegativity.

Nitrogen exists as diatomic nitrogen gas (N2) under standard conditions. The strong triple bond between nitrogen atoms accounts for its stability and inertness.

The strong triple bond in diatomic nitrogen (N2) requires a significant amount of energy (942 kJ mol-1) to break, making it relatively inert under normal conditions.

The absence of d orbitals and high electronegativity makes nitrogen different from other Group 15 elements. These factors contribute to its unique chemical behavior.