Chapter 2 Chemistry Structure and Bonding PDF

## 2.1 Structure and bonding ### States of matter and forces of attraction - Substances with high melting points have strong forces of attraction between their atoms (or ions). - Substances with low melting points have weak forces of attraction between their molecules. - **Solid:** strong bonding wi...

## 2.1 Structure and bonding ### States of matter and forces of attraction - Substances with high melting points have strong forces of attraction between their atoms (or ions). - Substances with low melting points have weak forces of attraction between their molecules. - **Solid:** strong bonding within molecules, strong attractive forces between molecules, particles are in fixed arrangements and do not move (they vibrate) - **Liquid:** strong bonding within molecules, weak forces between molecules, particles are close together and slide over each other. - **Gas:** weak forces between molecules, particles are far apart and move randomly. ### Covalent bonding - A covalent bond is formed by the force of attraction between the nuclei of two neighbouring atoms and a pair of electrons between them. - The attractive forces are in balance with the repulsive forces between the electron clouds when the nuclei are a certain distance apart. - Covalent bonds are usually strong and many energy is needed to break them. ### Sigma bonds and pi bonds - A covalent bond is formed when atomic orbitals overlap. - Each orbital which combines contributes one unpaired electron to the bond. - The "joined" orbital is called a molecular orbital. - The greater the overlap of the atomic orbitals, the stronger is the covalent bond. - **Sigma bonds (σ bonds):** formed by the overlap of atomic orbitals along a line drawn between the two nuclei. - **Pi bonds (π bonds):** formed by the sideways overlap of p atomic orbitals. ## 2.2 Covalent bonding - dot and cross diagrams ### Simple dot and cross diagrams - A shared pair of electrons in a single bond is represented by a single line. - It is usually only the electrons in the outer shell which are used in covalent bonding. - Electron pairs in the outer shell which are not used in bonding are called lone pairs. - **Dot and cross diagrams**: - Use a dot to represent the outer shell electrons from one atom and a cross to represent the outer shell electrons from another atom. - Draw the outer electrons in pairs to emphasise the number of bond pairs and the number of lone pairs. - If possible, electrons are arranged so that each atom has four pairs of electrons around it (an octet of electrons/ the noble gas electron configuration). Hydrogen is an exception - it can only have two electrons around its nucleus when forming covalent bonds. - When pairing electrons, a covalent bond is formed between an electron from one atom (dot) and an electron from another atom (cross). ### Electron deficient molecules - Some molecules are unable to complete the octet of electrons when they form covalent bonds. - These molecules are said to be electron deficient. - An example is boron trichloride, BCl3. This only has six electrons around the boron atom. ### Molecules with an expanded octet - Some molecules can increase the number of electrons in their outer shell to more than 8. - These molecules are said to have an expanded octet of electrons. - An example is sulphur hexafluoride, SF6. This has 12 electrons around the sulphur atom. ## 2.3 More dot and cross diagrams ### Molecules with multiple bonds - Some atoms form bonds by sharing two pairs of electrons. - A double bond is formed, represented by a double line, e.g. for oxygen, O=O - When atoms share three pairs of electrons, a triple bond is formed. - The triple bond is shown by a triple line, e.g. for nitrogen N=N. ### Co-ordinate bonding - A co-ordinate bond (dative covalent bond) is formed when one atom provides both the electrons for the covalent bond. - We need: 1. One atom with a lone pair of electrons. 2. A second atom with an unfilled orbital. - The atom with the unfilled orbital (an electron deficient atom) accepts the lone pair of electrons to complete the outer shell of both atoms. ## 2.4 Ionic and metallic bonding ### The formation of ionic compounds - When electrons are transferred from metal atoms to non-metal atoms, ions are formed. - The outer shell of both ions formed have the noble gas electron configuration. - Positive ions are formed by loss of electrons and negative ions are formed by gain of electrons. - In an ionic compound the number of positive and negative charges must balance. ### Dot and cross diagrams for ionic structures - The ions formed have a full outer shell of electrons (electron configuration of the nearest noble gas). - The charge on the ion is placed at the top right. - The square brackets indicate that the charge is spread throughout the ion. ### Ionic bonding - The ionic bond is an electrostatic force of attraction between oppositely charged ions. - The net attractive forces between these ions results in the formation of a giant ionic structure. - In this structure the ions are regularly arranged in a three-dimensional lattice. - The electrostatic attractive forces between the ions act in all directions and the bonding is very strong. ### Metallic bonding - Most metals exist in a lattice of ions surrounded a 'sea' of delocalised electrons. - Delocalised electrons are those which are not associated with any particular atom or ion. - They are free to move between the metal ions. - The number of delocalised electrons depends on the number of electrons lost by each metal atom. - The positive charges are held together by their strong electrostatic attraction to the delocalised electrons. - This strong electrostatic attraction acts in all directions. - So metallic bonding is usually strong - The strength of metallic bonding increases with: - increasing positive charge on the ions - decreasing size of the metal ions - increasing number of delocalised electrons. ## 2.5 Electronegativity and intermolecular forces ### Electronegativity - Electronegativity is the ability of a particular atom involved in covalent bond formation to attract the bonding pair of electrons to itself. - Electronegativity increases across a period from Group I to Group VII. - Electronegativity decreases down any group. - The order of electronegativity is: F > O > N > Cl > Br... > C > H. ### Polarity in molecules: bond polarisation - If the electronegativity values of the two atoms in a covalent bond are the same, we say that the bond is non-polar. - If the electronegativity values of the two atoms in a covalent bond are different, we say that the bond is polar. - In a non-polar molecule the centres of positive and negative charge coincide. - In a polar molecule the centres of positive and negative charge do not coincide. - The degree of polarity is measured by a dipole moment. ### Weak intermolecular forces - Intermolecular forces arise because of the attraction between the dipoles in neighbouring molecules.. - There are three types of intermolecular forces: - permanent dipole-dipole forces - van der Waals forces - hydrogen bonding. ### Hydrogen bonding - Hydrogen bonding is a special form of permanent dipole bonding. - It requires: 1. One molecule with an H atom covalently bonded to an F, O or N atom. 2. A second molecule having a F, O or N atom with a lone pair of electrons. ## 2.6 Properties of giant structures ### Lattices - A lattice is a regular three-dimensional arrangement of particles. - Lattices with strong bonding between the particles are called giant structures. - The three types of giant structures are: - giant ionic, e.g. sodium chloride, magnesium oxide - giant covalent (sometimes called giant molecular), e.g. diamond, silicon dioxide - metallic, e.g. copper, iron ### Properties of giant ionic structures - Melting and boiling points are high: it takes a lot of energy to break the large number of strong bonds between the oppositely charged ions. - Soluble in water: the ions can form ion-dipole bonds with water molecules. - Conduct electricity only when molten or dissolved in water: the ions can only when molten or dissolved in water. - Brittle: when a force is applied, the layers of ions slide. ### Properties of giant covalent structures - Melting and boiling points are high: it takes a lot of energy to break the large number of strong covalent bonds in the network of atoms. - Insoluble in water: the atoms are too strongly bonded to form bonds with water. - Apart from graphite they do not conduct electricity: there are no electrons or ions that are free to move. - Apart from graphite they are hard: the three-dimensional network of strong bonds in different directions is too difficult to break. ### Properties of metals - Melting and boiling points are high: it takes a lot of energy to break the large number of strong forces of attraction between the ions and the delocalised electrons. - Insoluble in water (although some react with water): the force of attraction between the ions and the delocalised electrons is too strong to form bonds with water - Conduct electricity when solid or molten: the delocalised electrons are free to move when a voltage is applied. - Malleable (can be beaten into shape) and ductile (can be drawn into wires): when a force is applied, the layers of metal ions can slide over each other. ## 2.7 Properties of simple molecular compounds ### Molecular lattices - Substances with a simple molecular structure such as iodine, I2, can form crystals. This reflects a regular packing of the molecules in a lattice. - The forces between the molecules are the weak van der Waals forces. - Simple molecular structures which are solids at room temperature are brittle. - They do not conduct electricity because they have neither delocalised electrons nor ions. ### Melting and boiling points of simple molecular structures - Many simple molecular substances are liquids or gases. - They have low melting and boiling points because it does not take much energy to overcome the weak intermolecular forces between the molecules. - Solids with molecular lattices, such as iodine or sulphur, are also easily broken down when heated. - van der Waals forces increase with: - increasing number of electrons in the molecule - increasing number of contact points in the molecule (contact points are places where the molecules come close together). ### The anomalous properties of water - Water has a much higher boiling point than expected by comparison with other Group VI hydrides. - The boiling points of the hydrides from H2S to H2Te increase gradually due to increased van der Waals forces. - Water's unusually high boiling point is due to hydrogen bonding, which is much stronger than other intermolecular forces. - Most solids are denser than their corresponding liquids. - Ice, however, is less dense than (liquid) water. This is because the molecules in ice are arranged in an 'open' lattice structure stabilised by hydrogen bonding. ### Solubility of simple molecular substances - In order to dissolve, the strength of the attractive forces between solute and solvent is usually greater than between the solute molecules themselves. - Most covalently-bonded molecules such as iodine, sulphur and butane, are insoluble in water. - Many are non-polar and so water molecules are not attracted to them. - They will, however, dissolve in non-polar solvents such as hexane if the new intermolecular forces formed between the solvent and solute are stronger than those between the solute molecules themselves. - Some simple molecules are soluble in water e a ethanol CH3CH2OH and ammonia NH3. This is because they can form hydrogen bonds with water. ## 2.8 Shapes of molecules (1) ### VSEPR theory - The valence shell electron pair repulsion theory (VSEPR theory) can be used to work out the shapes of molecules. - It uses the following rules: - Pairs of electrons in the outer shells of the atoms in a molecule repel each other and move as far apart as possible. This minimises repulsive forces in the molecule. - Repulsion between lone-pairs and lone-pairs of electrons is greater than the repulsion between lone-pairs and bond-pairs of electrons. - Repulsion between lone-pairs and bond-pairs of electrons) is greater than the repulsion between bond-pairs and bond-pairs of electrons. ## 2.9 Shapes of molecules (2) ### Hybridisation of orbitals - Methane has four C-H bonds of equal length. - The four unfilled C atomic orbitals can be thought of as being mixed so that each has *s* character and *p* character. - This process of mixing atomic orbitals is called hybridisation. - These mixed orbitals are called *sp³* orbitals - Hybridisation allows a greater overlap of atomic orbitals when a molecular orbital is formed and also allows each bond to be the same. - They are all σ bonds with equal repulsion between them. ### The structure of ethene - In ethene, one singly occupied 2*s* orbital and two of the three singly occupied 2*p* orbitals in each carbon atom hybridise to make three *sp²* orbitals. - These have similar shapes to *sp³* orbitals. - These *sp²* orbitals form σ bonds which are arranged in a plane making a bond angle of approximately 120° with each other since there is equal repulsion of the electrons. - The remaining 2*p* orbitals from each carbon atom overlap sideways to form a π bond. ### Resonance - In methane and ethene the electrons are localised, i.e. they are in particular positions. - In some substances, the molecular orbitals extend over three or more atoms, allowing some of the electrons free movement over these atoms. - These electrons are said to be delocalised. - Benzene, C6H6, has six carbon atoms arranged in a ring. - The bonds between the carbon atoms are neither double nor single bonds. They are somewhere in-between. - The composite structure is called a resonance hybrid. - In benzene, the six carbon atoms form a hexagon with three localised *sp²* hybrid orbitals (one to each hydrogen atom and two to other carbon atoms). - The 3 *sp²* orbitals are arranged in a plane. So the bond angles are 120°. - This leaves a single *p* orbital on each of the six carbon atoms. - These orbitals overlap sideways to form a delocalised system of π bonds. - The six electrons involved can move freely around the ring.