Chapter 4 - Structure and Properties of Substances

Chapter 4.1: Chemical Bonding

• Chemical bonds: electrostatic forces that hold atoms together in compounds lower-energy systems tend to have greater stability than higher-energy systems Bonded atoms are more stable
• Lewis structures: atomic symbol in the middle with its outer valence electrons around
  • Lines: bonding pairs between atoms
  • Dots: lone pairs
• Ionic bonding occurs with atoms of elements of large differences in electronegativity
  • s block metals paired with p block non-metals
  • Can be drawn with an orbital diagram or Lewis structures
  • Bonding allows atoms to obtain valence electron configuration of the closest Noble gas → octet rule is observed
• Ionic solids have:
  • Crystalline with smooth, shiny surfaces
  • Hard and brittle
  • Non-conductors of electricity and heat
  • High melting points
• Lattice energy: energy is given off when ionic crystal forms from the gaseous ions
  • The energy of the same amount is needed to break turn back to the gaseous ions
  • Adding energy decreases stability and processes that readily give off energy become stable
• Covalent bonding: electron-sharing to acquire noble gas configuration
• Bond energy: energy required to break the force of attraction between two atoms and to separate it
  • An increase of bonds means more bond energy
• Covalent compounds:
  • Can be soft solids, liquids, or gas at room temperature
  • Low melting points and boiling points
  • Poor conductors of electricity
  • Not usually soluble in water (acetone is an exception)
• Can use electronegativity difference to determine bonding type
• Ionic bonds: described as “electron transfer”, “losing electrons”, “gaining electrons”
• Polar covalent bond: covalent bond with unequal shared pair of electrons between atoms: ∆EN (change between electronegativities) between 0.4 and 1.7
• Bonded atoms with ∆EN between 1.7 and 3.3 are mostly ionic
• Metallic Bonding
  • Free-electron model: pictures metals as being composed of densely packed core of metallic cations with free roaming electrons
  • Metabolic bonding: force of attraction between positively charged cations and the pool of valence electrons that moves among them
  • Characteristics of metallic bonds:
  • Conductivity
  • Malleability and ductility
  • High melting and boiling points

Chapter 4.2: Molecular Shape and Polarity

• How to draw a Lewis Structure:
  • Least electronegative center (hydrogen and fluoride located at the edge)
  • Determine the total number of valence electrons of each atom (pay attention to charges)
  • Determine the total number of valence electrons to achieve the noble gas configuration
  • If a negative 2 change then add the total number of valence electrons to the calculated valence electrons
  • Add when negative charge, subtract when a positive charge
  • Subtract the number of valence electrons from a number of electrons needed to satisfy the octet rule
  • Divide it by two to give the number of bonds
  • Double bonds are two bonds and triple bonds are three bonds
  • Subtract the number of shared electrons from a number of valence electrons to get lone pairs to draw around the central atom
• Coordinate covalent bond: when a filled atomic orbital overlaps with an empty atomic orbital
• Resonance structure: models that give the same relative position of atoms but show different places for their bonding and lone pairs
  • Do not exit in reality, instead exists as a hybrid
• Expanded valence energy level: when a central atom has more than eight electrons in the valence energy level of the central atom
• Lewis models do not communicate about the molecule’s shape, you need 3D models for that
• Valence-Shell Electron-Pair Repulsion: bonding pairs and lone, non-bonding pairs of electrons in the valence level of an atom repel one another
  • LP-LP > LP-BP> BP-BP
  • LP (lone pair)
  • BP (bond pair)
  • Repulsion greatest in lone pair repulsions
• When all the electron groups are bonding pairs, a molecule will have one of the five geometrical arrangements. If one or more of the electron groups include a lone pair of electrons, variations in one of the five geometric arrangements result.
  • Linear: (2 BP)
  • Trigonal planar: (3BP) or (2 BP, 1 LP)
  • Tetrahedral: (4 BP) or (3 BP, 1LP) or (2 BP, 2 LP)
  • Trigonal bipyramidal: (5 BP) or (4 BP, 1LP) or (3 BP, 2 LP) or (2 BP, 3 LP)
  • Octahedral: (6 BP) or (5 BP, 1LP) or (4 BP, 2LP)
• How to find the molecular shape
  • Determine Lewis structure
  • Find the number of electrons around the central atom
  • Determine which five geometrical arrangement
  • Determine using positions of bonding pairs and lone pairs
• Dipole: charge separation for the entire molecule

Chapter 4.3: Intermolecular Forces in Liquids and Solids

• Intramolecular forces: forces exerted within a molecule or polyatomic ion
• Intermolecular forces: forces that influence physical properties of substances; forces of attraction and repulsion that act between molecules
• Dipole-dipole forces: attractions between opposite charges
  • Ex. Positive and negative
• Ion-dipole forces: the force of attraction between ion and polar molecule
  • Ex. Positive ion attracts with the negative end of the water
• Induced Intermolecular forces:
  • Ion-induced dipole forces: ion in close proximity to a nonpolar molecule that distorts electron density of non-polar molecule making it polarized
  • Ex. Nonpolar O2 with Fe2+
  • Dipole-induced dipole force: similar to an ion-induced dipole force, a polar molecule distorts the electron density of a nonpolar molecule making it polarized
• Dispersion (London) force: the force of attraction between nonpolar molecules larger number of electrons cause an uneven distribution of charge large shape of the molecule are stronger and results in increased boiling point
• Hydrogen Bonding: a strong form of dipole-dipole attraction that exists between a hydrogen atom in a polar-bonded molecule that contains bonds such as O, N, F. There are also unshared pairs of electrons from these electronegative atoms
  • Water pack in a specific way where the solid-state is less dense than the liquid state
  • Soluble for other polar covalent compounds such as alcohols
• Crystalline solids: organized particle arrangements of solids
  • Ex. Amethyst, garnet
  • Bonding and properties of Crystalline Solids:
    • Atomic solids: made up of individual atoms held by dispersion forces
    • Molecular solids: molecules held by dispersion forces, dipole-dipole forces, or hydrogen bonding
    • Network solids: continuous two or three-dimensional arrays
    • Carbon-based network solids:
    • Allotropes: different crystalline or molecular forms of the same element that differ in physical and chemical properties
• Amorphous solids: indistinct shapes because particle arrangements lack order
  • Ex. glass rubber
• Ionic Crystals: array of ions, arranged at regular positions in a crystal lattice
• The crystal lattice is made up of identical repeating unit cells which gives the crystal its unique shape

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