Please enable JavaScript.

Coggle requires JavaScript to display documents.

Structure and Bonding (Giant Ionic Structures (Properties of Giant Ionic…

- - - - Simple molecular structures have comparatively weak intermolecular forces of attraction, and boiling/melting a molecular substance will not require breaking the covalent bonds the molecules have, as the melting/boiling point size depends on the strength of the intermolecular forces.
        For example, if hydrogen bonds are present in the molecule, then the melting/boiling points of that substance will increase. Essentially, the greater the size of the molecule, there’s more van der Waals attractions formed, and to break these attractions/forces will require greater energy.
- - - - Dot and Cross diagram of Sodium Sulfide:
      - .
    - - .
- - - - Delta negative (δ-) represents the more electronegative ion, and delta positive (δ+) is used to represent a partial charge (less than unit charge associated with ions)
    - - Difference of electronegativity values (using Pauling scale of electronegativity)
        
        Nonpolar covalent = electronegativity difference = 0 to 0.4
        
        Polar covalent = electronegativity difference = 0.5 to 1.9
        
        Ionic = electronegativity difference = 2.0 to 4.0
  - - - Polar covalent
        
        Polar covalent: A chemical bond between two atoms with low electronegativity difference, in which the electrons required to form a bond is shared unequally between two atoms.
        The electrons from the bonded pair are attracted more towards the atom which is more electronegative (δ-).
        
        E.g: Water
        In a polar covalent bond, the electrons shared by the atoms are (on average) closer to the Oxygen nucleus than the Hydrogen nucleus.
        The electron pair is displaced toward the more electronegative atom.
        This Oxygen atom then obtains a partial-negative charge while the less electronegative atom, Hydrogen, has a partial-positive charge.
        
        Diagram of polar covalent bond:
      - Non-polar covalent
        
        Non-polar covalent: Non-polar covalent bonds are bonds between 2 non-metal atoms that have the same electronegativity values, and therefore, equal sharing of the bonding electron pair as the difference in electronegativities = 0.
        
        E.g: H2
        In H2, each H atom has an electronegativity value of 2.1.
        The difference between 2 H atoms in a H2 molecule is 0, therefore the covalent bond between them is considered nonpolar.
      - Polar molecule
        
        Polar molecule: A polar molecule is a molecule where the bond dipoles present do not cancel each other out, in which the bond dipoles present do not cancel each other out and thus results in a molecular dipole.
        Cancellation depends on the shape of the molecule and the orientation of the polar bonds.
        
        Molecular dipole
        
        A molecular dipole: Bond dipoles may/may not cancel out producing either molecules that are nonpolar (if they cancel out) or polar (if they do not cancel out)
        
        E.g: CO2 is a linear molecule with 2 bond dipoles that are equal and oppositely directed therefore the bond polarities cancel and the molecule is nonpolar.
        
        Diagram of carbon dioxide bond:
        
        1 more item...
        
        The molecules carbon dioxide, bromine fluoride and methane are all non-polar because the dipoles cancel out.
        
        Diagram:
        
        If the bonds are not symmetrically arranged, then the polarities will not cancel out.
        Any molecule with lone pairs will be polar unless the lone pairs cancel out with another lone pair.
        E.g: hydrogen chloride, chloromethane (CH3Cl) and ammonia (NH3)
      - Ionic bonds:
        
        Ionic bonds: Ionic bonding is the complete transfer of valence electron(s) between atoms, that generates two oppositely charged ions.
        The metal loses electrons to become a positively charged cation and the non-metal accepts those electrons to become a negatively charged anion.
        
        Diagram of NaCl ionic bonding:
      - Dative Covalent bonds:
        
        Dative covalent bond:
        A covalent bond is formed by two atoms sharing a pair of electrons, and held together because of the 2 positive nuclei attracting the electron pair.
        A dative covalent or co-ordinate bond is a covalent bond where it involves a shared pair of electrons, but both electrons come from the same atom.
        
        E.g: NH4
        N has a lone pair of electrons in NH3, and this lone pair of electrons (from one atom) will bond with a H+ to form NH4.
        
        Diagram of NH4 bonding:
- - - - A double covalent bond is formed where two pairs of electrons are shared between the atoms, rather than just one pair.
        Two oxygen atoms are joined together with two pairs of electrons shared between the atoms rather than just one pair, thus being a double bond.
- - - - The shape of the molecule is determined by the angles between the bonded atoms.
        The total number of charge centers around the central atom determines the geometrical arrangement of the electrons.
        
        Lone pairs of electrons have a higher concentration of charge than a bonding pair because the lone pairs are not shared between two atoms, so they cause more repulsion than bonding pairs.
        
        Order of repulsion
        
        Highest repulsion: Repulsion between two lone pairs.
        
        Repulsion between lone pair and bonding pair
        
        Lowest repulsion: Repulsion between two bonding pairs.
    - - 1) Molecules with two charge centres will position at 180° to each other (linear shape)
        
        E.d: CO2
      - 2) Molecules with 3 charge centres will position themselves at 120° to each other, (Planar triangular shape to electron distribution)
        
        E.g: BH3
      - 3) If there are lone pairs at the centre atom (rather than a bond), the angle will be less than 120°, to be 104.5° (Bent/V-shape)
        
        E.g: H2O
      - 4) Molecules with 4 atoms will position themselves at 109.5° to each other (tetrahedral shape to electron pairs)
        
        E.g: CH4
      - 5) If one or more of the charge centers are a non-bonding pair, the angle in which they position themselves will be less than 109.5°, to be 107° to each other. (Trigonal pyramidal)
        
        E.g: NH3
      - 6) Molecules with 5 charge centres will position them 120° to each other. (Trigonal bipyramidal)
        
        E.g: PCl5
      - 7) Molecules with six atoms are symmetrically arranged around the central atom, with 4 of the attachments positioned in a square plane with 90° bond angles.
        The remaining two attachments are positioned perpendicular to the square plane at opposite ends of the central atom.
        
        E.g: SF6
      - Summary of bonding:
      - 8) There are 2 lone pairs of electrons on opposite sides of the central atom. The other four atoms connect to the central atom. The bond angles are 90° to each other. (square planar)
        
        E.g: XeF4