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