Ionic and covalent Bonding
2.2.2 Bonding and Structure Metal atoms lose electrons to form +ve ions. Non-metal atoms gain electrons to form -ve ions. Mg goes from 1s2 2s2 2p63s2 to Mg2+ 1s2 2s2 2p6 O goes from 1s2 2s2 2p4 to O2- 1s2 2s2 2p6 Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges. E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl- ) Ionic Bonding The ions in an ionic solid are arranged in a regular 3D pattern called a giant ionic lattice Na+ ClThe sticks in this diagram are there to help show the arrangements of the ions. They do not represent the ionic bonds. Ionic bonding is between ions and all their surrounding oppositely charged ions. Each sodium ion in this structure is surrounded and equally attracted by six chloride ions. The ionic bond is the attraction between all these ions •High melting points – There are strong electrostatic attractive forces between the oppositely charged ions in the lattice •Non conductor of electricity when solid- The ions are held together tightly in the lattice and can not move so no charge is conducted •Good conductor of electricity when in solution or molten – The ions are free to move when in solution and molten. Charge can be carried Typical Physical properties of Ionic Compounds • They are usually soluble in aqueous solvents.Definition: An Ionic bond is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.Covalent Bonding Definition: covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms A Dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. A dative covalent bond is also called co-ordinate bonding. Common examples you should be able to draw that contain dative covalent bond (e.g. NH4+ , H3O+ , NH3BF3 ) O H H H + .. N H H H B Cl Cl Cl The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient Dative Covalent bonding The dative covalent bond acts like an ordinary covalent bond when thinking about shape so in NH4+ the shape is tetrahedral The term average bond enthalpy is used as a measurement of covalent bond strength. The larger the value of the average bond enthalpy, the stronger the covalent bond. Bonding Structure Examples Ionic : electrostatic force of attraction between oppositely charged ions Sodium chloride Magnesium oxide Covalent : shared pair of electrons Simple molecular: With intermolecular forces (Induced dipole–dipole, permanent dipole-dipole, hydrogen bonds) between molecules Iodine Ice Carbon dioxide Water Methane Giant Ionic Lattice Bonding and Structure. Only use the words molecules and intermolecular forces when talking about simple molecular substances Property Giant Ionic Molecular (simple) boiling and melting points high- because of giant lattice of ions with strong electrostatic forces between oppositely charged ions. low- because of weak intermolecular forces between molecules (specify type e.g Induced dipole–dipole/hydrogen bond) Solubility in water Generally good generally poor conductivity when solid poor: ions can’t move/ fixed in lattice poor: no ions to conduct and electrons are localised (fixed in place) conductivity when molten good: ions can move poor: no ions general description crystalline solids mostly gases and liquids
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2.2.2 Bonding and structure
Ionic bonding (a) ionic bonding as electrostatic attraction between positive and negative ions, and the construction of ‘dot-and-cross’ diagrams (b) explanation of the solid structures of giant ionic lattices, resulting from oppositely charged ions strongly attracted in all directions e.g. NaCl (c) explanation of the effect of structure and bonding on the physical properties of ionic compounds, including melting and boiling points, solubility and electrical conductivity in solid, liquid and aqueous states HSW1 Use of ideas about ionic bonding to explain macroscopic properties. Covalent bonding (d) covalent bond as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms (e) construction of ‘dot-and-cross’ diagrams of molecules and ions to describe: (i) single covalent bonding (ii) multiple covalent bonding (iii) dative covalent (coordinate) bonding ‘Dot-and-cross’ diagrams of up to six electron pairs (including lone pairs) surrounding a central atom. (f) use of the term average bond enthalpy as a measurement of covalent bond strength Learners should appreciate that the larger the value of the average bond enthalpy, the stronger the covalent bond. Definition and calculations not required. Average bond enthalpies and related calculations are covered in detail in 3.2.1 f.
Shape of molecules
Name No bonding pairs No lone pairs Diagram Bond angle Examples linear 2 0 180 CO2, CS2 , HCN, BeF2 Trigonal planar 3 0 120 BF3 , AlCl3 , SO3 , NO3 -, CO3 2- Tetrahedral 4 0 109.5 SiCl4 , SO4 2-, ClO4 -, NH4+ Trigonal pyramidal 3 1 107 NCl3 ,PF3 ,ClO3 ,H3O+ Bent 2 2 104.5 OCl2 , H2S, OF2 , SCl2 Octahedral 6 0 90 SF6 S. Shape of molecules. How to explain shape 1. State number of bonding pairs and lone pairs of electrons. 2. State that electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.) 3. If there are no lone pairs state that the electron pairs repel equally 4. If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs. 5. State actual shape and bond angle. Remember lone pairs repel more than bonding pairs and so reduce bond angles (by about 2.5o per lone pair in above examples)
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2.2.2 Bonding and structure
The shapes of simple molecules and ions (g) the shapes of, and bond angles in, molecules and ions with up to six electron pairs (including lone pairs) surrounding the central atom as predicted by electron pair repulsion, including the relative repulsive strengths of bonded pairs and lone pairs of electrons M4.1, M4.2 Learners should be able to draw 3-D diagrams to illustrate shapes of molecules and ions. HSW1,2 Using electron pair repulsion theory to predict molecular shapes. (h) electron pair repulsion to explain the following shapes of molecules and ions: linear, non-linear, trigonal planar, pyramidal, tetrahedral and octahedral Learners are expected to know that lone pairs repel more than bonded pairs and the bond angles for common examples of each shape including CH4 (109.5°), NH3 (107°) and H2O (104.5°).
Electronegativity and intermediate bonding
Electronegativity and intermediate bonding Definition Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself. F, O, N and Cl are the most electronegative atoms Factors affecting electronegativity Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more. It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent Formation of a permanent dipole – (polar covalent) bond A polar covalent bond forms when the elements in the bond have different electronegativities. When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends. The element with the larger electronegativity in a polar compound will be the δ- end H – Cl + – A compound containing elements of very different electronegativity and hence a very large electronegativity difference will be ionic e.g. CCl4 will be non-polar whereas CH3Cl will be polar N Goalby chemrevise.org 5 A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular ARE polar. Symmetric molecules The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no NET dipole moment: the molecule is NON POLAR C H H H Cl δ+ δ- CO2 is a symmetrical molecule and is a non-polar molecule
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2.2.2 Bonding and structure
Electronegativity and bond polarity (i) electronegativity as the ability of an atom to attract the bonding electrons in a covalent bond; interpretation of Pauling electronegativity values Learners should be aware that electronegativity increases towards F in the periodic table. HSW1,2 Using ideas about electronegativity to predict chemical bond type. (j) explanation of: (i) a polar bond and permanent dipole within molecules containing covalently-bonded atoms with different electronegativities (ii) a polar molecule and overall dipole in terms of permanent dipole(s) and molecular shape A polar molecule requires polar bonds with dipoles that do not cancel due to their direction. E.g. H2O and CO2 both have polar bonds but only H2O has an overall dipole.
Intermolecular bonding
Intermolecular bonding Induced dipole–dipole interactions Induced dipole–dipole interactions are also called London forces. They occur between all simple covalent molecules and the separate atoms in noble gases. In any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form. These temporary dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original one. Main factor affecting size of Induced dipole–dipole interactions The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Induced dipole–dipole interactions stronger between the molecules and so boiling points will be greater. Permanent dipole-dipole forces •Permanent dipole-dipole forces occurs between polar molecules •It is stronger than van der waals and so the compounds have higher boiling points •Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds) •Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms. The increasing boiling points of the alkane homologous series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the induced dipole–dipole interactions between molecules. Permanent dipole forces occur in addition to induced dipole–dipole interactions The increasing boiling points of the halogens down the group 7 series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the induced dipole–dipole interactions between the molecules. This is why I2 is a solid whereas Cl2 is a gas. The shape of the molecule can also have an effect on the size of the induced dipole–dipole interactions forces. Long chain alkanes have a larger surface area of contact between molecules for induced dipole–dipole interactions to form than compared to spherical shaped branched alkanes and so have induced dipole–dipole interactions. Induced dipole–dipole interactions occur between all molecular substances and noble gases. They do not occur in ionic substances. Permanent dipole–dipole and induced dipole–dipole interactions can both be referred to as van der Waals’ forces. The molecules are held further apart than in liquid water and this explains the lower density of ice O H H O H H O H H O H H O H H H2O H2S H2Se HF H2Te HCl HBr HI NH3 PH3 AsH3 SbH3 SiH4 CH4 GeH4 SnH4 100 200 300 400 Molecular mass 25 50 75 100 125 Boiling point K 7 N Goalby chemrevise.org It occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond. There is a large electronegativity difference between the H and the O,N,F Hydrogen bonding Always show the lone pair of electrons on the O,F,N and the dipoles and all the δ – δ + charges Hydrogen bonding occurs in addition to van der waals forces Water can form two hydrogen bonds per molecule, because oxygen is very electronegative, and it has two lone pairs of electrons. Molecular: Iodine There are covalent bonds between the Iodine atoms in the I2 molecule The crystals contain a regular arrangement of I2 molecules held together by weak induced dipole– dipole interactions intermolecular forces. Hydrogen bonding is stronger than the other two types of intermolecular bonding. The anomalously high boiling points of H2O, NH3 and HF are caused by the hydrogen bonding between the molecules The general increase in boiling point from H2S to H2Te is caused by increasing induced dipole– dipole interactions between molecules due to an increasing number of electrons. Alcohols, carboxylic acids, proteins, amides all can form hydrogen bonds
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2.2.2 Bonding and structure
Intermolecular forces (k) intermolecular forces based on permanent dipole–dipole interactions and induced dipole– dipole interactions Permanent dipole–dipole and induced dipole–dipole interactions can both be referred to as van der Waals’ forces. Induced dipole–dipole interactions can also be referred to as London (dispersion) forces. HSW1,2 Dipole interactions as a model to explain intermolecular bonding. (l) hydrogen bonding as intermolecular bonding between molecules containing N, O or F and the H atom of –NH, –OH or HF Including the role of lone pairs. (m) explanation of anomalous properties of H2O resulting from hydrogen bonding, e.g.: (i) the density of ice compared with water (ii) its relatively high melting and boiling points HSW1 Use of ideas about hydrogen bonding to explain macroscopic properties. (n) explanation of the solid structures of simple molecular lattices, as covalently bonded molecules attracted by intermolecular forces, e.g. I 2, ice (o) explanation of the effect of structure and bonding on the physical properties of covalent compounds with simple molecular lattice structures including melting and boiling points, solubility and electrical conductivity.
Credits: Neil Goalby