Periodic Table and Electronic Configuration

• Physical properties of atoms change dependent upon their position in the periodic table

The Periodic Table

• The periodic table is arranged by increasing atomic number
• Elements are grouped into periods (rows) and groups (columns)
• Vertical columns are called groups (e.g., Group 1 are alkali metals, Group 17 are halogens)
• Elements in the same group have similar properties and the same number of valence electrons
• Horizontal rows are called periods
• The number of valence electrons increases and properties change gradually across a period
• Sodium (Na) is in Group 1, Period 3 and has one valence electron, making it very reactive, like potassium (K)

Electronic Configuration

• Electron configuration is the arrangement of electrons in an atom's energy levels (shells) and subshells (s, p, d, f)
• Electrons fill the lowest energy levels first (Aufbau principle)
• Subshells include s (2 electrons), p (6 electrons), d (10 electrons), f (14 electrons)
• Focus on elements up to atomic number 36 (krypton)
• Notation is in the format 1s² 2s² 2p⁶, etc.
• Sodium (Na, atomic number 11) has an electron configuration of 1s² 2s² 2p⁶ 3s¹
• Chlorine (Cl, atomic number 17) has an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁵

Periodicity

• Periodic trends are how properties like atomic radius, ionic radius, ionization energy, and electronegativity change across the periodic table
• Atomic radius decreases across a period (left to right) as nuclear charge pulls electrons closer
• Atomic radius increases down a group as more electron shells are added
• Ionization Energy (IE) is the energy it takes to remove an electron
• Ionization Energy (IE) increases across a period because of a stronger attraction
• Ionization Energy (IE) decreases down a group because electrons are farther from the nucleus
• Electronegativity measures the ability to attract electrons in a bond
• Electronegativity increases across a period and decreases down a group
• Fluorine (F) has a higher ionization energy and electronegativity than sodium (Na) because F is further to the right in its period

Reactivity Trends

• Reactivity depends on how easily an element gains or loses electrons
• Reactivity increases down a group for metals (e.g., K is more reactive than Na) because the valence electron is farther from the nucleus so it's easier to lose
• Reactivity increases up a group for nonmetals (e.g., F is more reactive than Cl) because smaller atoms attract electrons more strongly
• Group 1 metals react vigorously with water, forming hydroxides (e.g., 2Na + 2H₂O → 2NaOH + H₂)

Ionic Lattice Structures

• Ionic compounds form when metals lose electrons to non-metals, creating positive (cations) and negative (anions) ions
• Ions arrange into a 3D lattice held by electrostatic forces
• Properties of ionic compounds include high melting and boiling points because of strong ionic bonds
• Ionic compounds are soluble in water because ions dissociate
• Ionic compounds conduct electricity when molten or dissolved because ions are free to move
• Ionic compounds are brittle as the lattice shatters if layers are misaligned
• Sodium chloride (NaCl) is an example
• Sodium (Na) loses 1 electron to form Na⁺ and chlorine (Cl) gains 1 electron to form Cl⁻
• Sodium chloride forms a cubic lattice structure and dissolves in water: NaCl(s) → Na⁺(aq) + Cl⁻(aq)

Properties of Metals

• Metals have a lattice of positive ions that are surrounded by a "sea" of delocalized electrons, also known as metallic bonding
• The properties of metals include being good conductors of heat and electricity because delocalized electrons are able to move
• The properties of metals include being malleable and ductile because layers of ions can slide without breaking the bond
• The properties of metals include high melting points because of strong metallic bonds
• Copper (Cu) is used in wires because it's a good conductor and ductile

Metallic Bond Strength

• The strength of metallic bonds depends on the number of delocalized electrons and the charge of the metal ions
• Bond strength increases along with more valence electrons and a higher ion charge
• Magnesium (Mg, 2 valence electrons) has stronger metallic bonds than sodium (Na, 1 valence electron), thus Mg has a higher melting point

Electronegativity (EN) and Bond Polarity

• Electronegativity (EN) measures an atom's ability to attract electrons in a bond
• Bonds are polar where atoms in a bond have different EN values (unequal sharing of electrons)
• Non-polar bonds have an EN difference= 0 (e.g., Cl₂)
• Polar bonds have an EN difference of 0.4-1.7 (e.g., HCl, EN difference = 0.9)
• Ionic bonds have an EN difference greater than 1.7 (e.g., NaCl, EN difference = 2.1)
• In H₂O, O (EN = 3.5) is more electronegative than H (EN = 2.1), so the O-H bonds are polar

Molecular Polarity

• Molecular polarity depends on bond polarity and its shape
• Symmetrical molecules are non-polar if the bonds are polar but the dipoles cancel out (e.g., CO₂, which is linear and non-polar)
• Asymmetrical molecules are polar if dipoles do not cancel out (e.g., H₂O, which is bent and polar)
• NH₃ (trigonal pyramidal) is polar because the dipoles do not cancel out

Intermolecular Forces (IMFs)

• IMFs are forces between molecules and are weaker than covalent bonds
• Types of IMFs include London Dispersion Forces (LDF), temporary dipoles in all molecules, and stronger with larger molecules
• Types of IMFs include Dipole-Dipole Forces, which occur between polar molecules
• Types of IMFs include Hydrogen Bonding, which is a strong dipole-dipole bond between H and N, O, or F (e.g., in H₂O)
• Stronger IMFs lead to higher boiling points
• H₂O has a higher boiling point (100°C) than H₂S (no hydrogen bonding, -60°C)

Properties of Covalent Molecules

• Properties depend on the type of covalent structure
• Simple molecules have low melting/boiling points (weak IMFs), and don't conduct electricity (e.g., I₂)
• Giant covalent networks have high melting points, are hard, and don't conduct (except graphite), e.g., diamond (tetrahedral) and graphite (layers)
• Diamond is hard and does not conduct electricity
• Graphite conducts electricity from delocalized electrons

Chromatography

• Chromatography separates mixtures based on how substances interact with a stationary phase (e.g., paper) and a mobile phase (e.g., solvent)
• Substances with stronger IMFs to the stationary phase move slower
• Polar dyes move slower in a water solvent during paper chromatography of inks

Review: Functional Groups

• Functional groups determine the properties of organic molecules
• Examples of functional groups:
  • OH (alcohol, e.g., ethanol, C₂H₅OH)
  • COOH (carboxylic acid, e.g., ethanoic acid, CH₃COOH)
  • NH₂ (amine, e.g., methylamine, CH₃NH₂)

Properties Within a Homologous Series

• A homologous series is a family of organic compounds with the same functional group and general formula (e.g., alkanes: CₙH₂ₙ₊₂)
• As the chain length increases:
  • Boiling point increases (stronger LDF)
  • Solubility in water decreases (less polar)
• Methane (CH₄) has a lower boiling point (-161°C) than pentane (C₅H₁₂, 36°C)

Structures and Allotropes: C and Si

• Allotropes are different forms of the same element with different structures
• Carbon Allotropes:
  • Diamond: Tetrahedral, hard, doesn't conduct electricity
  • Graphite: Layered, conducts electricity (delocalized electrons)
  • Graphene: Single layer of graphite, very strong
• Silicon: Similar to diamond (tetrahedral) and has weaker bonds, used in semiconductors

Plastics and Their Structure

• Polymers are large molecules made of repeating monomer units
• Types of polymers:
  • Addition Polymers: Formed by adding monomers with double bonds (e.g., polyethene from ethene, CH₂=CH₂)
  • Condensation Polymers: Formed by monomers reacting with the loss of a small molecule (e.g., nylon, polyester)
• Properties depend on structure (e.g., cross-linking makes plastics stronger)
• Polyethene (addition polymer) is flexible and used in plastic bags

Bonding is a Continuum

• Bonding types (ionic, covalent, metallic) are not absolute, they exist on a spectrum
• The bonding triangle plots bonds based on electronegativity differences and metallic character
  • Ionic: High EN difference (e.g., NaCl)
  • Covalent: Low EN difference (e.g., Cl₂)
  • Metallic: Delocalized electrons (e.g., Cu)
• AlCl₃ has some covalent character despite being mostly ionic (EN difference = 1.5)

Across the Bonding Triangle

• Properties change across the bonding triangle
• Ionic bonds have a high melting point, are brittle, and conduct electricity when molten
• Covalent bonds have a low melting point (molecular) or very high melting point (giant covalent) and do not conduct electricity
• Metallic bonds have a high melting point, are malleable, and conduct electricity
• SiO₂ (giant covalent) has a high melting point (1610°C) and is hard, unlike NaCl (ionic, brittle)