CAPE Chemistry: Period 3 Elements

• Module 3 of CAPE CHEMISTRY covers topics related to Period 3 elements (Sodium to Argon), Group II elements, Group IV elements, Group VII elements, and Transition elements.

Period 3: Sodium to Argon Trends

• Atomic radius decreases across Period 3 (Na to Cl).
• Nuclear charge increases across Period 3.
• The number of shells remains constant, and the attraction between the nucleus and outermost electrons increases leading to the decrease in atomic radii.
• Ionic radius generally decreases across Period 3.
• Metallic character changes to non-metallic character, with the radius of the metallic ion being smaller than its corresponding atom.
• Radius of the non-metallic ion is larger than its corresponding atom.
• From Na to Al, ionic radius decreases as the ratio of protons to electrons increases, increasing nuclear attraction on the remaining electrons.
• From P to Cl, ionic radius decreases as additional electrons increase electron repulsion.
• Melting Point increases from Na to Si and then decreases from P to Ar.
• Na, Mg, and Al are metals with a giant metallic structure where ions are held in a sea of delocalized electrons.
• From Na to Al, nuclear charge and the number of delocalized electrons increase, resulting in an increase in metallic bond strength.
• Si has a giant molecular structure with strong covalent bonds and requires large amounts of energy to break its lattice structure in melting.
• P, S, and Cl exist as small, simple molecules held together by weak Van der Waals forces, requiring increasingly less energy to melt as the molecule gets smaller, except for sulfur, which has a higher melting point than phosphorus due to its larger size and stronger Van der Waals forces.
• Electrical conductivity increases from Na to Al and then decreases from Si to Ar.
• Na, Mg, and Al are metals containing delocalized electrons capable of carrying charge and conductivity increases from Na to Al as the number of delocalized electrons increases.
• Conductivity decreases at Si due to its metalloid nature and tightly held covalent bonds, only allowing for a few electrons to become delocalized at high temperatures.
• Phosphorus, Sulfur, Chlorine, and Argon do not conduct electricity because their electrons are held in covalent bonds and are not free to move.
• Argon exists as a single atom with outermost electrons that are not free to move due to being held in a stable energy level.
• Electronegativity, the ability of an atom to attract electrons in a covalent bond, increases across Period 3.
• This increase is because of increasing nuclear charge and decreasing atomic radius, with a negligible shielding effect, which increases nuclear attraction for electrons.
• Density increases from Na to Si and then decreases from P to Ar.
• From Na to Si, the forces of attraction between atoms increases causing them to be more closely packed, hence increasing density.
• P, S, and Cl are simple molecules, held together by weak Van der Waals forces that are not closely packed, and their density decreases.

Reactions of Period 3 Elements

• Reactions with Oxygen:
• Na burns vigorously to give Na₂O and Na₂O₂.
• Mg burns with a brilliant flame, forming MgO.
• Al burns to form a layer of Al₂O₃.
• Si burns at high temperatures to form SiO₂.
• P burns vigorously, forming P₄O₆ and P₄O₁₀ (Phosphorus V oxide).
• S burns to form SO₂ and SO₃.
• Cl and Ar have no reactions.
• Reactions with Chlorine:
• Na burns in chlorine with a bright orange flame to form NaCl.
• Mg burns with a white flame to form MgCl₂.
• Al burns to form AlCl₃.
• Si reacts with powder to form SiCl₄.
• P burns in chlorine to produce PCl₃ and PCl₅.
• S reacts to form S₂Cl₂.
• Cl and Ar have no reaction.
• Reactions with Water:
• Na reacts violently to form NaOH and H₂.
• Mg reacts with water to form Mg(OH)₂, and reacts with steam to form MgO and H₂.
• Al slowly reacts with steam and powdered Al to form Al₂O₃ and H₂.
• Si, P, and S have no reaction, as they are insoluble.
• Cl reacts to form HCl and HClO.
• Ar has no reaction.

Oxidation State Variation

• The highest oxidation state of an element in Period 3 equals the number of electrons in the outermost shell, meaning all valence electrons are involved in bonding.
• Oxides and Chlorides Oxidation States:
• Na₂O: +1 state
• MgO: +2 state
• Al₂O₃: +3 state
• SiO₂: +4 state
• P₄O₁₀: +5 state
• SO₃: +6 state
• Cl₂O₇: +7 state
• NaCl: +1 state
• MgCl₂: +2 state
• AlCl₃: +3 state
• SiCl₄: +4 state
• PCl₅: +5 state

Period 3 Oxides

• Sodium Oxide (Na₂O):
• Bonding: Ionic
• Melting/Boiling Point: High
• Acid/Base Nature: Basic
• pH: 13
• Reacts with H₂O to form NaOH, is soluble and reacts to form NaOH.
• Magnesium Oxide (MgO):
• Bonding: Ionic
• Melting/Boiling Point: High
• Acid/Base Nature: Basic
• pH: 11 -Reacts with H₂O to form Mg(OH)₂, is slightly soluble.
• Aluminum Oxide (Al₂O₃):
• Bonding: Ionic with covalent character
• Melting/Boiling Point: High
• Acid/Base Nature: Amphoteric
• Reacts with H₂O to form salt and water.
• Silicon Dioxide (SiO₂):
• Bonding: Giant covalent
• Melting/Boiling Point: High
• Acid/Base Nature: Acidic
• Does not react with H₂O.
• Phosphorus Pentoxide (P₄O₁₀):
• Bonding: Simple covalent
• Melting/Boiling Point: Low
• Acid/Base Nature: Acidic
• pH: 3
• Vigorous Rreaction with H₂O to form phosphoric acid.
• Sulfur Trioxide (SO₃):
• Bonding: Simple covalent
• Melting/Boiling Point: Low
• Acid/Base Nature: Acidic
• pH: 2
• Vigorous Reaction with H₂O to form sulfuric acid.
• Dichlorine Heptoxide (Cl₂O₇):
• Bonding: Simple covalent
• Melting/Boiling Point: Low
• Acid/Base Nature: Acidic
• pH: 2
• Vigorous Reaction with H₂O to form acid.

Period 3 Chlorides

• Sodium Chloride (NaCl):
• Bonding: Ionic (held together by oppositely charged ions)
• Dissolves in water, High Melting/Boiling Point, Neutral pH (7)
• Magnesium Chloride (MgCl₂):
• Bonding: Ionic
• Dissolves in water, High Melting/Boiling Point, Neutral pH (6.7)
• Anhydrous - Aluminum Chloride (Al₂Cl₆):
• Bonding: Ionic with covalent character,, reacts vigorously with water to form white HCl fumes.
• Low Melting/Boiling Point, Acidic pH (3)
• Silicon Tetrachloride (SiCl₄):
• Bonding: Simple Covalent, reacts vigorously with water to form acid.
• Low Melting/Boiling Point, Acidic pH (2)
• Chlorine (Cl₂):
• Bonding: Simple Covalent, reacts with H₂O to form HCl and HOCl, Low Melting/Boiling Point, Acidic pH (2)
• Phosphorus Pentachloride (PCl₅):
• Bonding: Simple Covalent, "Reacts with H₂O forming acid and HCl fumes"., Low Melting/Boiling Point, Acidic pH (2)
• Sulfur Chloride (S₂Cl₂):
• Bonding: Simple Covalent, "forming HCl fumes, SO₂ and S.", Low Melting/Boiling Point, Acidic pH (2)

Acid/Base Nature

• Acid/Base Nature of Period 3 Oxides:
• Metal Oxides (Na & Mg) - Basic
• Aluminum Oxide - Amphoteric
• Non-metal Oxides - Acidic
• Acid/Base Nature of Period 3 Metal Hydroxides:
• Na/Mg - Basic
• Al - Amphoteric

Uses of Period 3 Compounds

• Aluminum Hydroxide:
• Used in antacid medication as it reacts with excess acid in the stomach, preventing acid reflux and heartburn.
• Phosphorus:
• Used in flares to produce billowing smoke to draw attention to rescuers.
• Red phosphorus is used on matchboxes as friction converts it to white phosphorus, which ignites spontaneously in air.
• Argon:
• Used in fluorescent and incandescent lighting to stop oxygen from corroding the filament.

Group II Elements Trends

• Atomic Radius: increases down Group II because there is one more shell, causing the screening effect to increase.
• Melting Point: Decreases down the group is generally true with the exception of magnesium.
• Ionization Energy: First Ionization Energy decreases down the group because there is less effective nuclear force.

Reactions of group II elements

• The E values for Group II increase in negativity down the group.
• Since reverse reactions are all positive, eEelectrons can be easily lost from these elements.
• Barium loses electrons easiest, and reactivity increases down group II.
• Reaction with Oxygen: All Group II elements react with oxygen to form their metallic oxide, and the vigor of the reaction increases down the group.
• Reaction with Water: Beryllium does not react with water.
• Magnesium reacts very slowly with liquid water, but rapidly with steam to form magnesium oxide and hydrogen gas.
• Calcium, Strontium, and Barium all react with cold water, increasing in vigor to give the metal hydroxide and hydrogen gas.
• These hydroxides are not very soluble, but solubility increases down the group, so less precipitate is observed moving down the group.
• Reaction with acids: All Group II elements react with acids to produce salts and hydrogen gas with the vigor of reactions increasing down the group.
• Solubility of Group II Sulfates: decreases down the group

Reactions with water

• Beryllium (Be) does not react with water.
• Magnesium reacts very slowly with cold liquid water, but rapidly with steam.
• Calcium, Strontium and Barium all react with cold water, increasing in vigour to give the metal hydroxide and hydrogen gas.
• These hydroxides are not very soluble, But, soloubility increases down the group, so less precipitate is observed moving down the group.
• All group II elements react with acids to produce salts and hydrogen gas. Magnesium + 2HCl -> MgCl2 + H2

Group II Solubility

• The solubility of group II sulfates decreases down the group.
• Dissolving a group II sulfate encompasses three processes: dissociation (reverse of lattice energy), enthalpy of hydration of cation Mg2+ and enthalpy of hydration of anion: SO4.
• The size of the cation increases down the group, while the lattice energy and hydration energy of group two elements decreases down the group, thereby decreasing the solubility of sulfates formed and hence increasing the amount of precipitate.
• Comment on the solubility of the sulfate of Radium in water: insoluble (since solubility decreases).

Thermal Decomposition Compounds

• The nitrates of group II elements undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen.
• More stable than nitrates down the group, so higher temperatures are required to decompose them.
• Group II nitrate stability increases down the group.
• Smaller ions polarize the nitrate more easily to form the oxide compounds:

CABONATES

• The carbonates of group II elements undergo thermal decomposition to give the metal oxide and carbon dioxide.
• As you move down the group, the element becomes more stable, so more energy is required when moving down..
• The smaller ions at the top of the small group will be more stable than larger ones at the bottom.
• NOTE*
• The first reaction is used to treat acidity.
• The second reaction is what allows us to counter act acidity in soils.
• The last reaction is used in water treatment to reduce acidity and is used as a dietary supplement.

Group IV elements

• ATOMIC RADIUS
• Moving down group IV, the atomic radius increases.
• As a result of successive elements having one more shell, the screening effect also increases as full inner shells shield the outer electrons from increasing nuclear charge.
• The distance between the nucleus and outermost electron increases.
• Metallic CHARACTER
• Metallic character increases down the group covalent to giant metallic structures.
• ELECTRICAL CONDUCTIVITY
• Graphite is a conductor, while silicon and germanium are semi-conductors. -Tin and Lead ARE METALS!!!!

GROUP 4 Compounds

• Molecules within group 4 elements all form tetrachlorides with the formula XCl4.
• They are simple covalent molecules with a tetrahedral shape.
• Group IV elements form the centre of the molecule with 4 covalent bonds to chlorine atoms.

Group 4 Notes

• All tetrachlorides have a have low melting/boiling point.
• The molecules aee volatile at room temperature, and the x-cl bond gets longer down the group and the molecule gets weaker.
• As one moves down the group the Waals forces become weaker, and renders the molecule less stable.
• Ail tetrachlorides with the exception of CCly (tetrachloromethane-which is immiscible) are readily hydrolysed by water to the oxide in the tu oxidation state, producing acidic fumes of HCT. The ease of hydrolysis increases down the group.

Oxides of Group IV

• CO2 and SiO2 are stable even at high temperatures.
• ONLY Significant decomposition from PbO2 to PbO upon heating.
• CO is readily oxidized to CO2, it has a strong triple bond and does not decompose on heating.
• Silicon monoxide (SiO) is readily oxidized to silicon dioxide (SiO2).
• Tin oxide SnO and lead oxide (PbO), does not decompose on heating in presence of air.
• All the compounds in Oxide structure need to have Ionic with covalent character.

Group IV Bond Info

• All elements in Group IV have four electrons in their valence shell, Going down the group, Ge, S, and Pb tend to form sonic compounds which have a +2 oxidation state. In these compounds: Get, Sn, and Pb ions are formed. When two electrons are lost from the valence shell The two remising S electrons on the valence shell and relatively westable stable and not easily removed.
• All +2 state are said to be weakly acidic/Neutral

Oxidation Characteristics Of Groups IV

• The group exists with oxidation numbers +2, -4, and -5
• The +4, oxidation state being most common
• The +2 oxidation state becomes more stable as you move down the line.
• The +4 oxidation state is less stable moving down the group and undergoes thermal decomposition

Elements

• Carbon and Silica are very stable
• Geranium and Tin are less stable
• Lead and Lead Dioxide are least stable.

Uses of Ceramics

• Furnace linings made of Heaters and Boilers.
• Abrasives.
• Manufacture of Glass and Porcelain

Group VII Elements

• As one descends the graph in Atomic Radius, the atom increases as each element has an additional shell, so a shielding effect will occur which reduces effective nuclear charge.
• Volatility decreases
• Density increases down the group
• Colors tend to darken as you move down the columns.

State

• Down the group there increase in Vander waals force from increasing in state.
• All agents are oxydising agents

Actions of group VII Elements

• Halogens are oxydising compounds

Additional Notes on Halogens

• Sodium Thiosulphate is used to look at Iodine and Bleach Sample. Sodium Thiosulphate's can all be feasible with the other 7th group compounds.

Halogen reactions

• Hydrogen Halides have a reaction that happens directly with Halogens (The reaction goes slower moving down the group. Hydrogen Flouride is highly reactive. Hydrogen Chloride is also highly reactive (In the presence of Sunlight).
• Silver Nitrate can be dissolved with Ammonia, and will display the Rxn.
• Hydride Stability can occur down the chain of elements due to loss in Electrons.
• The strength of Halides (when used to make molecules) will increase has you descend the chain of elements

Reactivity

• Halogens react with different types of compounds.
• Cold Sodium Hydroxide and chloride form well.
• Chloride 6 at zero with two bonds to NaOh makes : the same setup with 1 negative connection to Nano, and one positive
• In these reactions are called disproportionation
• They are reactions in which a single substance becomes oxidized and Reduced.

Notes On Transition Elements

• A transition element is an element which forms ions with a partially. filled subshell.
• Scandium and zinc not considered metals They form compounds with variable oxidation states. They form cloured compounds. They form complex pons + They are good catalysts. They exhibit paramagnetisim.