covalent bonds and forces

• Properties of non-metallic substances

  Types of bonds
• Intramolecular bonds


  The strong bonds between the carbon and oxygen atoms within carbon dioxide molecules
• Intermolecular bonds


  The weak bonds between carbon dioxide molecules
• Weak intermolecular bonds


  The bonds that are broken when molecular substances such as water boil or melt, allowing the molecules to separate from each other while the atoms within the molecules remain bound to one another
• Covalent bonding


  • Covalent bonded molecules are formed from a non-metal and a non-metal
  • A covalent bond is formed by the sharing of electrons
  • Due to their atomic structure, non-metallic atoms like to share electrons in order to complete their valence shell
  • Most of the time, this is due to the octet rule which describes that atoms are most stable when they have eight electrons in their valence shell
  • The electrons that are shared between the atoms are attracted to the positive nucleus of both atoms, and this electrostatic force of attractions holds the atoms together
  • Non-metals have high electronegativities so they can attract electrons easily but don't give up electrons easily
• Single bonds


  When a bond is formed by the sharing of two electrons, it is referred to as a single covalent bond
• Double bonds


  It is also possible for an atom to share more than one electron, for example oxygen which has six valence electrons needs two electrons to complete its outer shell, so a double covalent bond can form by the sharing of a total of four electrons (or two pairs of electrons)
• Triple bonds


  Some atoms have the ability to share three electrons, for example nitrogen which has five valence electrons requires three more electrons to fill its valence shell, so when two nitrogen atoms bond together, each atom shares three electrons with the other, forming a triple covalent bond
• Electron Dot Diagrams


  A form of shorthand that simplify how we can draw the outer shell electrons in atoms, also known as Lewis diagrams
• Structural Formula


  Used to show the covalent bonds which are represented as solid lines, with the lone pairs (non-bonding electrons) absent
• Valence structure


  Displays the bonding electrons as solid lines and the non-bonding electrons (lone pairs) as dots
• Bonding Electrons are electrons shared between the atoms
• Non-Bonding Electrons


  The outer shell electrons not involved in forming a bond, known as lone pairs, which are important in determining the shape of the molecule
• Valence shell electron pair repulsion theory (VSEPR)

  Provides a way of predicting what the shapes of molecules will be according to the repulsion of electron pairs, as like charges repel each other and negatively charged pairs of electrons are most stable when they are as far apart as possible
• Tetrahedral shape


  4 bonding pairs of electrons and 0 lone pairs around the central atom, with the four pairs of bonding electrons positioned as far apart as possible in a tetrahedral arrangement
• Trigonal pyramidal shape


  3 bonding pairs of electrons and 1 lone pair, with the three hydrogen atoms forming a pyramidal arrangement with the nitrogen atom
• 1. shaped or bent


  2 bonding pairs of electrons and 2 lone pairs, with the two hydrogen atoms forming a V-shape or bent arrangement with the oxygen atom
• Linear shape


  1 bonding pair of electrons and 3 lone pairs, with the hydrogen atom and fluorine atom forming a linear molecule
• Double and triple bonds


  • If a central atom has two single bonds and one double bond, the three bonds will repel each other resulting in a trigonal planar shape
  • If a central atom has a single bond and a triple bond, the molecule is linear
  • If a central atom has two double bonds, the two double bonds repel each other resulting in a linear molecule with a bond angle of 180°
• Electronegativity


  A measure of how strongly an atom can attract electrons to itself, with fluorine being the most electronegative atom
• Polarity


  In the context of chemistry, means that there is a partial positive and partial negative charge present
• Non-Polar bonds


  Formed when atoms of the same or similar electronegativity are covalently bonded together, with the electron pairs shared equally between the two atoms
• Polar bonds


  Formed when two atoms with a relatively large difference in electronegativities are covalently bonded together, with the shared electrons spending more time in the vicinity of the atom with higher electronegativity
• Partial charges


  When two atoms bond that have different electronegativities, the atom with the higher electronegativity has a greater share of the electron pair and carries a partial negative charge, while the other atom carries a partial positive charge
• Polar Molecules


  Contain polar bonds and are not symmetrical in their shape, with an unequal distribution of electrons forming a slightly positive and slightly negative end of the molecule
• Polar Molecule Examples


  • Asymmetrical
• Non-Polar Molecules


  Contain polar or non-polar bonds and are symmetrical in their shape, with the partial charges cancelling each other out
• Non-Polar Molecules Examples


  • Symmetrical
• Non-Polar molecule with Polar bonds

  • Both water and carbon dioxide have polar covalent bonds, but carbon dioxide is linear and symmetrical, so the partial charges cancel each other out making it a non-polar molecule
• Non-Polar molecule with Polar bonds


  • Both CCl4 and CHCl3 have polar covalent bonds, but CCl4 is symmetrical so the partial charges cancel each other out making it non-polar, while CHCl3 is polar as it is non-symmetrical