Bonding

• Ionic bonding

  The transfer of electrons from one element to another, creating ions with opposite charges. Between these ions there is an electrostatic attraction. The attraction holds the ions together in a lattice. Ionic bonding happens between metals and non-metals
• Metallic bonding

  Has a lattice of positive metal ions within a 'sea' of delocalised outer electrons. Whilst the positive ions repel each other, they are held together by their electrostatic attraction to the electrons. This attraction gives metals fairly high melting and boiling points. Metals are good conductors of heat and electricity because of their mobile electrons. They are strong, malleable, and ductile
• Covalent bonding is the sharing of electrons
• Giant covalent (macromolecular) compounds such as silicon dioxide, diamond, and graphite, are large lattices of covalent bonds. The bonds are very strong, so require large amounts of energy to break. Therefore, such compounds have high melting and boiling points. Silicon dioxide and diamond do not conduct electricity but graphite does
• Ionic lattices will not conduct electricity when solid, but when they are molten or dissolved the ions are free to move and so they can carry charge
• Dative covalent bonding

  Occurs when both electrons in a covalent bond come from the same atom. A dative bond is represented using an arrow instead of a line. It has exactly the same properties as any other covalent bond
• Properties of diamond and graphite

  • Diamond: 4 C-C bonds at each atom, lattice structure, hard, does not conduct electricity.
  • Graphite: only 3 C-C bonds instead of 4, soft layers can slide over each other, delocalised electrons between layers that carry charge
• Pauling scale

  A measure of electronegativity: the more electronegative an atom (higher number), the more it will attract electrons
• The more electronegative an atom, the more it will attract electrons
• Electronegativity

  The power of an atom to attract the electron density in a covalent bond towards itself
• Down a group

  Electronegativity decreases because the shielding increases and the distance between the nucleus and outer electrons also increases
• Electronegativity

  The power of an atom to attract a pair of electrons in a covalent bond
• Across a period

  Electronegativity increases because the nuclear charge increases (increase number of protons) while the number of shells remains the same
• Electronegativity

  It depends on:

  • The nuclear charge (number of protons)
  • The distance between the outer electron and the nucleus
  • The amount of shielding by shells
• Forces between molecules

  The three intermolecular forces:

  • Permanent dipole-dipole forces
  • Van Der Waais ( induced dipole-dipole)
  • Hydrogen bonding
• Strength of metallic bonds (factors that affect it)

  • Charges since the greater the charge of the positive metal ion the strong metallic bond to electrons
  • The size of ion, the larger the atomic radius means there's more shielding which causes a weaker attraction between ion and delocalised electrons so a weaker bond
• Metallic bond

  Its the electrostatic attraction between positive metal ions and delocalised electrons
• Why are metals malleable and ductile ?

  Metals have layers that allow it to slide
• Giant ionic lattices properties

  • Solid at room temperature because it has strong bonds
  • High boiling and melting point because it has strong ionic bonds which means a strong electrostatic attraction from the opposing charged ions.
  • When dissolved in water or molten they carry a charge
  • It's brittle since the layers slide so ions with the same negative or positive charge would come in contact and repel from each other.
• VSEPR theory

  Lone pairs repel more than bond pairs

  So they create the shape of the molecule
• Van Der Waais forces

  • The strength of Van Der Waais forces increased as molecular size increases
  • Atomic radius increases down a group so the outer electron is further away from the nucleus. So there's a weaker electrostatic attraction between the electron and nucleus and so a temporary dipole are easier to induce
• Permanent dipole- permanent dipole forces

  • Permanent dipole- dipole forces arise if there is a difference in electronegativity between atoms in a bond and that those bond are unequally distributed in a molecule
• Hydrogen bonds

  • Hydrogen bonds can only happen to nitrogen, oxygen and fluorine
  • Because the atoms are highly electronegative due to their high nuclear charge and small atomic radius. This cause some electrons to be withdrawn from the hydrogen atom causing it to be slightly positive
  • It's the strongest intermolecular forces
• Melting and boiling points in period 3 elements

  • For Na, Mg and Al they all have high melting and boiling points due to strong metallic bonds. So a lot of energy is needed to break the electrostatic attraction between sea of delocalised electrons and positive metal ion
  • Si is a macromolecular molecule so many strong covalent bonds between atoms so a lot of energy is needed to break the bonds. Leading it to have a high melting and boiling point.
• Melting and boiling point of period 3 elements pt2

  • Cl₂(g), S₈(s) and P4(s) are simple molecular so they have weak Van der Waais between molecules so little energy is needed to break them however S₈ has a higher boiling and melting point then P4 because it has more electrons which means stronger van Der Waais forces.
  • Ar is monoatomic so it has very weak Van Der Waais between atoms
• Factors affecting the strength of ionic bond

  • Charge of ions
  • Size of ions
• Why are ionic crystals brittle


  The layers in an ionic crystals slide causing ions with the same charge to be touching which would make them repel from each so the crystals would snap
• Factors that affect the soulubilty of ionic compounds

  • Strength of ionic bond
  • Strength of attraction between ions and water