S3.1 The Periodic Table : Classification of Elements

S3.1 The Periodic Table : Classification of Elements

S3.1.1 and S3.1.2 Periods, Groups and Blocks/ Periodicity and Electronic Configuration:

⭐️ Periodicity is the regular repetition of properties of elements arising from patterns in their electronic configuration.

• Elements are placed in ascending order of atomic number (Z)
• Vertical columns form groups
• Horizontal rows form periods
• Elements in the same group have the same number of outer shell electrons
• Elements in the same period have the same highest main energy level
• Metallic elements found on the left (s block)
• Non metals found on the right (p block)
• Some central metals/transition elements (d block)
• Lanthanoids and actinoids found at the bottom (f block)
• Alkali metals – G1
• Halogens – G17
• Noble gases – G18
• Metalloid elements have properties of both metals and non metals
  • Physical properties resemble metals
  • Chemically, they are similar to non metals

S3.1.3 Periodicity in Properties of Elements :

📌Atomic/Ionic Radius :

• Atomic radius is half of the internuclear distance between neighbouring nuclei
• It increases down a group due to more electron shells, which increases the shielding effect
• It decreases across a period as the addition of protons increases the effective nuclear charge experienced by an element’s valence electrons – pulling them toward the nucleus
• Ionic radius increases down a group
• Ionic radii of positive ions < atomic radii (parent atom) – due to loss of an outer energy level
• Ionic radii of negative ions > atomic radii (parent atom) – addition of electrons into outer energy level causes increase in repulsions, and electrons move further apart
• Across a period – atomic radius of positive ions decreases with increase in charge
• Across a period – atomic radius of negative ions decrease as charge decreases

📌 Ionization Energies :

• Ionization energy is the energy required to remove the outermost electron from a gaseous atom
• Represented by the equation : M(g) –> M+(g) + e^–
• Ionization energies decrease down a group as electron shielding increases and it is easier to remove valence electrons
• Ionization energies increase across a period as there are more protons and hence a higher effective nuclear charge is experienced by valence electrons
• Departures from these trends serve as evidence for subshells

📌Electron Affinity :

• The first electron affinity of an element is the energy change that takes place when 1 mole of electrons is added to 1 mol of gaseous atoms to form 1 mol of gaseous ions
• Represented by the equation : X(g) + e^– –> X^–(g)
• First electron affinity is usually exothermic – electron goes from infinity to experiencing nuclear attraction
• Magnitude of electron affinity decreases down a group
• Across period – more exothermic

📌Electronegativity :

• Electronegativity is a measure of the attraction of an atom in a molecule for the electron pair in the covalent bond of which it is part
• Decreases down a group – bonding electrons are furthest from nucleus
• Increases across a period – increase in nuclear charge causes increased attraction between nucleus and bonding electrons

S3.1.4 Periodicity in Reactivity :

• Increasing metallic character of group 1 elements
• Decreasing non metallic character of group 17 elements
• Group 18 : Noble Gases
  • Unreactive
  • Other elements try to achieve noble gas configuration
• Group 1 : Alkali Metals
  • Highly reactive – form M⁺ ions
  • Metallic bonding (inc in metallic character)
  • Reactivity inc down a group
  • Good conductors of heat and electricity
  • Shiny grey surfaces when cut with a knife
  • Reaction with oxygen : M⁺(s) + O₂(g) –> 2 M₂O (s) [basic oxide]
  • Reaction with water : 2 M(s) + 2 H₂O (l) –> 2 MOH(aq) + H₂(g) [alkaline solution]
• Group 17 : Halogens
  • Non metals, diatomic molecules
  • Reactivity decreases down a group (weak F-F bond in F₂)
  • Relative reactivity can be seen through series of displacement reactions
    • Cl₂(aq) + 2 KBr (aq) — 2 KCl (aq) + Br₂ (aq)
    • Cl₂(aq) + 2 KI (aq) –> 2 KCl (aq) + I₂(aq)
    • Br₂(aq) + 2 KI(aq) —> 2 Kbr(aq) + I₂(aq)
  • Halogens form insoluble salts with silver – called halides
    • Ag⁺(aq) + X^–(aq) –> AgX (s)

S3.1.5 Metal and Non-metal Oxides :

• Metallic oxides – basic
• Non metallic oxides -acidic
• Oxides of metals from Na to Al – form giant ionic structures
• Oxides of P, S, Cl – molecular covalent structures
• Oxide or Si – giant covalent
• Ionic character of compound depends on difference in electronegativity between both elements in that compound
• Oxides become more ionic down a group – as difference in electronegativity increases
• Oxides only conduct electricity in aqueous or molten solutions where ions are free to move

Examiners Tip : Note that the maximum oxidation state for a period 3 element is related to its group number. Eg +1 for elements in G1.

• Basic oxides react with acid to form salt and water
• General Equation for G1 oxide with water :
  • M₂O (s) + H₂O (l) –> 2 MOH (aq)
• General Equation for G2 oxide with water :
  • MO (s) + H₂O (l) –> M(OH)₂(aq)
• Amphoteric oxides react with acids and bases
• Beryllium/Aluminium form amphoteric oxides
  • Al₂O₃(s) + 6H⁺(aq) –> 2 Al³⁺(aq) + 3 H₂O(l)
  • Al₂O₃(s)+ 2OH^–(aq) + 3 H₂O (l) –> 2 Al(OH)₄^–(aq)
• Acidic oxides react with bases to form salt
  • SO₂(g) + H₂O (l) –> H₂SO₃(aq)
• SO₃ and CO₂ are also acidic oxides

📌 Ocean Acidification :

⭐️Ocean acidification is the reduction of pH in oceans over an extended period of time due to increased CO₂ uptake from the atmosphere

• Equilibrium set up by 4 equations between CO₂ in atmosphere and CO₂ in water
  • CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq)
  • H₂CO₃(aq) ⇌ HCO₃^–(aq) + H⁺(aq)
  • HCO₃^–(aq) ⇌ CO₃^2-(aq) + H⁺(aq)
• Rainwater is naturally acidic due to dissolved CO₂
• Acid rain – rain with a pH lower than 5.6
• Sulfuric and nitric oxides lower pH
• Acid deposition – acidic substances deposited on surface of Earth
• Sulfuric oxides produce acid rain
  • S(s) + O₂(g) –> SO₂(g)
  • 2 SO₂(g) + O₂(g) —> 2 SO₃(g)
  • SO₃ (g) + H₂O (l) –> H₂SO₄(aq)
• Nitric oxides produce acid rain
  • N₂(g) + O₂(g) –> 2 NO(g)
  • 2 NO (g) + O₂ (g) –> 2 NO₂(g)
  • NO₂(g) + HO* (g) –> HNO₃(g)

🔍TOK Connect : Although all rainwater is acidic, the term ‘acid rain’ only refers to some water. How does language influence communication and dissemination of knowledge in science?

🌍 Real World Perspective : Acid rain affects buildings and structures leading to corrosion, especially structures of limestone. It also affects marine ecosystems, leading to dissolution of calcium carbonate skeletons in reefs and crustaceans. Acid rain water accumulates higher levels of toxic metals like Cadmium and Aluminium, linked to diseases like kidney dysfunction and Alzheimer’s.

S 3.1.6 Oxidation States :

• Oxidation state is the charge an atom would have if the compound was composed of ions
• Oxidation is
  • loss of electrons
  • loss of hydrogen
  • addition of oxygen
  • increase in oxidation state
• Reduction is
  • gain of electrons
  • gain of hydrogen
  • decrease in oxidation state

📌 Rules to Assign Oxidation States :

1. If the compound is ionic, oxidation state is charge on the ion.
2. Atoms in uncombined elements have an oxidation state of 0.
3. Oxidation states of all atoms in a neutral compound must add up to 0.
4. Oxidation state of Oxygen is -2 except for in peroxides, where it is -1.
5. Oxidation state of H is +1 except in metal hydrides where it is -1.
6. Sum of oxidation states of atoms in a compound should be equal to the charge on the coumpound
7. Maximum oxidation state = number of electrons in outer shell. For transition metals – sum of electrons in s and d subshells.

• Roman numerals used for oxidation states in names of compounds
• Oxyanion – negative ions with element + oxygen

	IUPAC Name	Trivial Name
SO42-	Sulfate (VI) ion	Sulfate ion
SO32-	Sulfate (V) ion	Sulfite ion
NO3–	Nitrate (V) ion	Nitrate ion
NO2–	Nitrate (III) ion	Nitrite ion

S3.1.7 Discontinuities in Patterns of First Ionization Energy [HL] :

• Exceptions to general trend of increasing ionisation energy : Boron’s first ionization energy is lesser than Berylliums
  • B – 1s²2s²2p¹
  • Be – 1s²2s²
  • p orbital higher in energy – easier to remove
• Oxygen’s first ionization energy is lesser than Nitrogen’s
  • O – 1s²2s²2p⁴
  • N – 1s²2s²2p³
  • In Oxygen – electrons are paired up, greater repulsion and easier to remove

S3.1.8/9 Characteristic Properties of Transition Elements/ Variable Oxidation States [HL]:

⭐️Transition elements have partially filled d subshells or form positive ions with a partially filled d subshell

• Due to small increase in effective nuclear charge, atomic radii starts decreasing
• This similarity in atomic radii is important to understand the ability of d block metals to form alloys – atoms can be replaced without disruption of solid structure
• Physical Properties
  • High melting point
  • High denistry
  • High electrical and thermal conductivity
  • Malleable
  • Ductile
• Chemical Properties
  • Exhibit more than one oxidation state in compounds and complexes
  • Form complex ions
  • Coloured complexes
  • Act as catalysts in many reactions
    • Eg. Iron in the Haber process of Nickel in the conversion of alkenes to alkanes
  • Can exhibit magnetic properties due to presence of unpaired electrons in d orbitals
    • Iron, Nickel and Cobalt show strong magnetic properties

🌍Real World Perspective : The properties and characteristics of transition metals make them important commodities internationally. The mining and extraction of these elements is important for the economic development of many countries.

🧠 Examiners Tip : Remember Zinc is not a transition metal as it does not have an incomplete d subshell.

• All transition metals show +2 and +3 oxidation states
• Oxidation states above +3 generally show covalent character
• Compounds with higher oxidation states tend to be oxidizing agents

3.1.10 Coloured Complexes [HL] :

• Transition metals in solution have a high charge density and attract water molecules to form coordination bonds with the positive ions to form a complex ion
• Ligand is a negative/neutral molecule with a lone pair of electrons (lewis base)
• Coordination number is the number of coordination bonds from the ligand to the central metal ion
• Transition elements except Titanium/Scandium form an octahedral complex ion of the form [M(H₂O)₆]²⁺
• The color of transition metal ions is related to the presence of partially filled d orbitals
• The color of a substance is determined by which colors of light (wavelengths of light) it absorbs and hence reflects
• Octahedral complex ion
  • When energy is absorbed to promtoe electrons, the d subshells split into 2 groups, 3 higher and 2 lower in energy
  • Whatever color of light is absorbed, the complementary color is reflected and seen

🧠 Exam Tip : A partially filled d-shell is required for a complex ion to be colored. Sc³⁺/Ti⁴⁺ have no electrons in 3d and Cu⁺/Zn²⁺ have 10 3d electrons.

📌 Factors that affect the color of Transition Metal Complexes :

• Identity of metal
  • Nuclear charge – greater repulsion, greater splitting
• Oxidation State of metal
  • Electronic configuration of ions differs
  • For complex ions containing same metal and ligands, greater oxidation state = greater splitting
• Coordination number and shape
  • different shapes – changes arrangement of splitting
• Nature of ligand
  • Spectrochemical series : I^–<Br^–<Cl^–<F^–<OH^–<H₂O<NH₃<CO< CN^–
  • Based on how much they cause d orbital to split

• Absorbance of a compound at a fixed wavelength is directly proportional to its concentration
• Calibration curve measures absorbance at a select wavelength over a range of known concentrations
• Absorbance of solution to be analysed is measured at the same wavelength as standard solutions and concentration of unknown solution can be determined using a calibration curve