1.4 - bonding

• Boiling temperature


  The temperature at which a substance changes from a liquid state to a gaseous state
• Bond angle


  The angle that is found between two bonds from the same atom in a covalently bonded compound
• Coordinate bond


  A type of covalent bond in which one bonding atom provides both electrons in the bonding pair
• Covalent bond


  The strong electrostatic attraction between two nuclei and the shared pair of electrons between them
• Electron pair repulsion


  Pairs of electrons around a nucleus repel each other so the shape that a molecule adopts has the pairs of electrons positioned as far apart as possible
• Electronegativity


  The ability of an atom to attract the bonding electrons in a covalent bond. The most electronegative elements (N,O,F) are small and have a relatively high nuclear charge
• Electrostatic attraction


  The attraction between 2 species with opposite charges
• Hydrogen bonding


  An interaction between a hydrogen atom and an electronegative atom, commonly nitrogen, fluorine or oxygen. The slightly positive hydrogen is attracted to the lone pair on the electronegative atom. Hydrogen bonds are stronger than van der Waals and dipole-dipole forces but weaker than ionic and covalent bonds
• Intermolecular forces


  The forces which exist between molecules. The strength of the intermolecular forces impact physical properties like boiling/melting point
• Ion


  An ion is formed when an atom/molecule loses or gains electrons. This gives it an overall charge - a positive charge if it has lost at least one electron and a negative charge if it has gained at least one electron
• Ionic bond


  Strong electrostatic attraction between two oppositely charged ions. The strength of attraction depends on the relative sizes and charges of ions
• Ionic compound


  A compound made up of anions and cations which are held together by ionic bonds, which arise due to the electrostatic attraction between oppositely charged ions. These structures are neutral overall
• Linear


  The shape of a molecule when the central atom has 2 bonding pairs and no lone pairs of electrons. The bond angle is 180°
• Octahedral


  The shape of a molecule in which the central atom has 6 bonding pairs. The bond angle is 90°
• Permanent dipole-dipole forces


  When molecules with polar covalent bonds interact with dipoles in other molecules dipole-dipole intermolecular forces are produced between the molecules. These intermolecular forces are generally stronger than van der Waals forces but weaker than hydrogen bonding
• Polar bond


  A covalent bond between two atoms in which the electrons in the bond are unevenly distributed. This causes a slight charge difference, inducing a dipole in the molecule
• Solubility


  The ability of a given substance to dissolve in a solvent
• Solvent


  A liquid that can dissolve other substances
• Tetrahedral


  The shape of a molecule when the central atom has 4 bonding pairs. The bond angle is 109.5°
• Trigonal planar


  The shape of a molecule when the central atom has 3 bonding pairs and no lone pairs of electrons. The bond angle is 120°
• Van der Waals


  Also known as induced dipole–dipole, dispersion and London forces, van der Waals forces exist between all molecules. They arise due to fluctuations of electron density within a nonpolar molecule. These fluctuations may temporarily cause an uneven electron distribution, producing an instantaneous dipole. This dipole can induce a dipole in another molecule, and so on
• Ionic bonding


  1. Metal atom loses electrons and becomes a cation (+)

  2. Non-metal gains electrons and becomes an anion (-)

  3. Strong electrostatic forces of attraction between oppositely charged ions
• Dot and cross diagrams


  Show how ionic bonds form
• Covalent bonding


  1. Non-metals bond together, each atom gives one electron to form a bond pair with opposed electron spins

  2. Electrostatic attraction between positive nuclei and shared electrons
• Dot and cross diagrams


  Show how covalent bonds form
• Single bonds


  • Water (H2O)
• Double bonds


  • Oxygen (O2)
• Coordinate bonding


  Same as a covalent bond but both electrons forming the bond pair come from the same atom
• Very few compounds are purely ionic or purely covalent, many have bonds that are intermediate between pure covalent and pure ionic</b>
• Polar bonds


  Bonds that are intermediate in character between pure covalent and pure ionic, degree depends on difference in electronegativity
• The greater the difference in electronegativity, the greater the ionic character
• Electronegativity


  Measure of how strongly atoms attract electrons in a covalent bond, higher value means better element can attract bonding electrons
• Polar covalent bond


  Bonding electrons are pulled towards the more electronegative atom, which takes up a slightly negative charge
• Non-polar covalent bond


  Atoms have equal electronegativities, electrons are equally shared
• Electronegativity difference ranges


  • Less than 0.4 - non-polar covalent
  • 0.4 to 1.9 - polar covalent
  • 2.0 or more - ionic
• Dipole-dipole forces


  Attraction between slightly positive and slightly negative regions of polar molecules
• Induced dipole-induced dipole forces


  Temporary dipoles in non-polar molecules induce dipoles in neighbouring molecules, creating attraction
• Van der Waals forces


  Dipole-dipole and induced dipole-induced dipole forces
• Hydrogen bonds


  Special intermolecular forces between hydrogen atoms bonded to small highly electronegative elements with lone pairs
• Hydrogen bonds


  • Increase boiling points
  • Increase solubility