Bonding and Shapes of Molecules

• Definition: a chemical bond between two ions of opposite charge.
• Between a metal and a non-metal.
• A metal will lose an electron and become positive and a non-metal will gain this electron and become negatve. The opposite charges attract and the two ions are held together by an electrostatic attraction.

• Definition: a bond in which two atoms share one or more pairs of electrons. A hydrogen molecule H-H has a single covalent bond, where one pair of electrons is shared. An oxygen molecule O=O has a double covalent bond, where two pairs of electrons are shared.
• Most group 4 elements make double bonds as it requires too much energy to lose 4 electrons. 
• In a covalent bond, the positive nuclei of both atoms attract the negative charge of the bonding pair of electrons 
  • Bonding pair: a shareed pair of electrons in a covalent bond.
• Covalent molecules are neutral.
• They can either exist as simple covalent molecules or as a giant covalent structure.
• In a covalent bond, an atom donates an electron. In a dative bond, both electrons in a single covalent bond come from one atom.
  • Dative covalent bond: a covalent bond in which the shared electron pair originates from the same atom. It can be written as X-->Y, showing that the shared electron pair originated from X.

• Definition: the bonding found in metals. The positive metal ions are attracted to the delocalised electrons and this electrostatic attraction holds the structure together.
• In solid state, metals are neither ionic or covalent.
• Each metal ion transfers its outer electrons to become a positive ion. The electrons make a sea of delocalised electrons which are free to move throughout the metal. The electrostatic attraction holds the metal together.

• Properties of a substance depend on its bonding. The attraction between one molecule and another is intermolecular bonding.
• Physical properties of a solid depend on arrangement of particles and the forces between them. 
• A structure may be crystalline or amorphous.
• In a crystalline structure, the particles are arranged in a fixed repeating pattern. This pattern gives crystals their shape .
• In an amorphous structure, there is no regular pattern to the way particles are arranged.
• There are four main types:
  • Ionic
  • Metallic
  • Macromolecular
  • Molecular 

• Definition: crystals consisting of positively and negatively charged ions held together in a regular arrangement by their electrostatic charges.
  • Key ionic crystal: sodium chloride.
• Ionic crystals consist of a lattice of positive and negative ions. 
  • Each Na+ is surrounded by six Cl- ions, and each Cl- is surrounded by six Na+ ions.
• An ionic crystal melts when there is enough heat energy transferred to overcome strong electrostatic attractions. When this happens, the ions move from their fixed place and are free to move. 
• Electrostatic attractions are strong so so the melting point of an ionic crystal is high
  • Sodium chloride's melting point is 801oC.
• To conduct electricity, the compound must have ions free to carry electrical charge.
  • Sodium chloride does not conduct electricity in solid form as ions are held in a fixed position. But it is free to conduct when molten 

• Definition:crystals consisting of positive metal ions that are held together by electrostatic attraction between the ions and a sea of delocalised electrons.
  • Key metallic crystal: magnesium.
• Metallic crystals consist of positive metal ions and a sea of delocalised electrons.
  • Each Mg2+ ion has tweleve adjacent ions, six on the same plane, three above and three below. This is called hexagonal close packed. This basic pattern is repeatd, making small crystals or grains.
• A metallic crystal melts when there is enough heat energy transferred to overcome strong electrostatic attractions between the positive ions and the delocalised electrons
• Metallic bonds are strong, so much energy is needed to break these bonds, so metals usually have high melting points.
  • Magnesium's melting point is 650oC.
• To conduct electricity, a compound must have free ions/electrons to carry electrical charge.
  • Magnesium does conduct electrcity as the delocalised electrons are free to carry charge.

• Also called giant covalent crystals.
• Definition: a crystalline solid in which its atoms are linked together by covalent bonds.
  • Key macromolecular crystals: diamond, graphite.
• Group 4 elements can form macromolecular crystals. Diamond and graphite are allotropes of carbon - they have different structures but are in the same state.
  • In diamond each carbon is covalently bonded to four other carbons giving a tetrahedral structure. This regular symmetrical pattern makes diamond the hardes naturally occuring substance. Diamond has a very high melting point - 3500oC. There are no charged particals, so diamond does not conduct.
  • In graphite each carbon is covalently bonded to three other carbons in hexagonal rings, in flat sheets. The fourth carbon makes up a cloud of delocalised electrons that hold the structure togetherwith VDW.Graphite has a very high melting point - 3230oC. Graphite can conduct electricity because of its delocalised electrons that are free to move.

• Definition: a solid formed when molecules are bonded together in a lattice pattern that repeats throughout the crystal
  • Key molecular crystal: iodine, ice.
• Molecular crystals have molecules held together by weak intermolecular forces 
  • I2 is covalent but it forms crystals with a regular arrangement of molecules. In iodine, molecules are held together by VDW. When iodine melts, the intermolecular forces break and I2 is free to move. This is because VDW are weak. Iodine has a low melting point - 113.5oC. No charged particles so cannot conduct.
  • When water freezes, ice crystals form. Weak hydrogen bonds form between water molecules. The water molecules are held in a regular pattern. When ice melts intermolecular forces break and water molecules can move. THE COVALENT BONDS DO NOT BREAK. Intermolecular forces are weak, so ice has a low meltion point - 0oC. Ice is not charged so it does not conduct.

• Shapes are important when finding explanations for the properties of a substance. The shape is determined by the number of elecrons pairs around the central atom. Electron pairs repel each other and are arranged as far apart as possible. This determines the shape of the molecule. 
• In a molecule, atoms have two different types of electron pairs:
  • Bonding pair: a shared pair of electrons in a covalent bond
  • Lone pair:a pair of electrons not involved in bonding.
• Lone pairs have a greater repulsion than bonding pairs because they are closer to the nucleus than the bonding pair. A lone pair reduces the angle by 2.5o.

See table.

• In a covalent bond between to hydrogen atoms the electrons are distributed evenly. The electron density is symmetrical because each hydrogen has the same power of attraction.
  • Electron density: a measure of an electron occurring in a specific location in an atom or molecule.
• In a covalent bond between a hydrogen and a fluorine, the electrons are drawn closer to the fluorine. The electron density is unsymmetrical because the fluorine attracts the electrons more stronly than the hydrogen. This means the negative charge from the electrons is closer to the fluorine, so the fluorine is slightly more negative than the hydrogen.
  • Electronegativity: the tendency of an atom to gain electons. Elements whose atoms gain electrons easily are the most electronegative.
• The bond in hydrogen fluoride is a polar covalent bond. This makes the molecule polar with a pernament dipole. This is not always the case. 
• What effects electronegativity? - size of the atom (distance between nucleus and outer electron), shielding and nuclear charge. 

• Types of dipole: permanent dipole, instantaneous dipole (electrons move constantly an at one moment the electrons may not be evenly distributed, so one end of a molecule may be slightly negative. This only lasts briefly, so is called instantaneous). A dipole may be induced by a non-polar molecule being close to a polar molecule.
• If a molecule has a dipole, they will attract; there are three types: per-per, per-ind, ind-ind. Ind-ind is also known as Van der Waals forces. These are very weak intermolecular forces. 
• Hydrogen bonding forms between a hydrogen and an oxygen, fluorine or nitrogen. Hydrogen bonding is weak, but stronger than VDW.