Ionic Bonding and Ionic Radii

Ionic Bonding

• Metal atoms lose electrons to become positive ions.
• Non-metal atoms gain electrons to become negative ions.
• Magnesium (Mg) becomes Mg2+ by losing electrons: 1s2 2s2 2p63s2 to 1s2 2s2 2p6
• Oxygen (O) becomes O2- by gaining electrons: 1s2 2s2 2p4 to 1s2 2s2 2p6
• Ionic crystals form giant lattices.
• Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions formed by electron transfer.
• Stronger ionic bonding and higher melting points occur when ions are smaller and/or have higher charges, e.g., MgO has a higher melting point than NaCl.

Ionic Radii

• Positive ions are smaller than their corresponding atoms because they have one less electron shell.
• The ratio of protons to electrons increases, resulting in a greater net force holding the remaining electrons more closely.
• Negative ions from groups five to seven are larger than their corresponding atoms.
• Negative ions have more electrons than protons, reducing the attraction per electron and making the ion bigger.
• N3-, O2-, F-, Na+, Mg2+, and Al3+ all have the same electronic structure as the noble gas Neon (Ne).
• Moving from N to F and then Na to Al, the number of protons increases while the number of electrons remains the same.
• Effective nuclear attraction per electron increases, and ions get smaller.
• Within a group, ionic radii increase as you go down because the ions have more electron shells.

Evidence for the Existence of Ions

• Electron density maps from X-ray diffraction show the likelihood of finding electrons in a region.
• Contours on the maps represent lines of equal electron density.
• The maps demonstrate that, for NaCl, the ions are arranged in a regular pattern, and chloride ions are larger than sodium ions.
• Ions are discrete because electron density falls to zero between them.
• Measuring ion radius from electron density maps is difficult because the edge of the ion cannot be precisely determined.

Physical Properties of Ionic Compounds

• High melting points are due to strong attractive forces between the ions.
• They are non-conductors of electricity in solid form because the ions are held tightly and cannot move.
• They are conductors of electricity when dissolved in solution or molten, as ions are free to move.
• They are brittle and easy to cleave because a small force will push ions with similar charges next to each other, the repulsion forces these apart.

Migration of Ions

• Colored ions migrate in an electric field: the blue color of Cu2+ ions moves to the negative electrode, and the yellow of CrO42- ions moves to the positive electrode.
• Potassium manganate (MnO4-), which is purple, migrates to the positive electrode when a voltage is applied.

Covalent Bonding

• A covalent bond is strong and is caused by the electrostatic attraction between the bonding shared pair of electrons and the two nuclei.
• The strength of covalent bonds can be demonstrated by the high melting points of giant atomic structures.
• High melting points occur since they contain many strong covalent bonds in a macromolecular structure.
• X-ray diffraction for the hydrogen molecule shows high concentration of negative charge between Hydrogen nuclei strongly attracted by both positive nucleui.
• In a covalent compound there is significant electron density between the atoms.
• Nuclei joined by multiple bonds have a greater electron density between them.
• This causes a greater force of attraction between the nuclei and the electrons between them, resulting in a shorter bond length and greater bond strength.

Dative Covalent Bonding

• 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 a coordinate bond.
• Common examples include NH4+, H3O+, and NH3BF3.
• In the ammonium ion (NH4+), the dative covalent bond acts like an ordinary covalent bond, so the shape is tetrahedral.
• Two aluminum chloride (AlCl3) molecules join together through two dative bonds to form the dimer Al2Cl6.
• The arrow direction goes from the atom providing the lone pair to the electron deficient atom.

Molecular Shapes

• Molecular shape is determined by the number of bonding pairs and lone pairs of electrons.
• Electron pairs repel each other and try to get as far apart as possible which produces a position of minimum repulsion.
• If there are no lone pairs, the electron pairs repel equally.
• Lone pairs repel more than bonding pairs.
• The shape and bond angle are then determined and will depend on the number of bonding pairs and lone pairs of electrons present.
• Lone pairs repel more than bonding pairs so reduce bond angles by about 2.5° per lone pair.
• Some basic shapes include linear, trigonal planar, tetrahedral, trigonal pyramidal, and bent.
• More complex shapes are seen in variations of octahedral and trigonal bipyramidal where some bonds are replaced with lone pairs.

Electronegativity

• Electronegativity is the relative tendency of an atom in a covalent bond to attract electrons in that bond to itself.
• Electronegativity is measured on the Pauling scale (ranges from 0 to 4).
• Fluorine, oxygen, nitrogen, and chlorine are the most electronegative atoms.
• Fluorine is the most electronegative element and has a value of 4.0.
• Electronegativity increases across a period as the number of protons increases and the atomic radius decreases.
• It decreases down a group as the distance between the nucleus and the outer electrons increases, and shielding increases.
• Ionic and covalent bonding are extremes of a continuum; electronegativity differences determine where a compound lies on this scale.
• Compounds with similar electronegativity have small electronegativity differences (< 1.7) and are purely covalent.
• Compounds with very different electronegativity have large electronegativity differences (> 1.7) and will be ionic.
• A polar covalent bond occurs when elements in the bond have different electronegativities (around 0.3 to 1.7).
• A polar covalent bond has an unequal distribution of electrons, producing charge separation (dipole) with δ+ and δ- ends.
• The element with larger electronegativity will be the δ- end.
• Symmetric molecules (all bonds identical and no lone pairs) will not be polar, even if individual bonds are polar, because the individual dipoles 'cancel out', and there is no net dipole moment.
• CCl4 is non-polar, whereas CH3Cl is polar.

Intermolecular Forces

• London forces occur between all molecular substances and noble gases but not in ionic substances.
• London forces are also called instantaneous, induced dipole-dipole interactions.
• In any molecule, electron density fluctuates.
• This can cause parts of the molecule to become more or less negative, forming temporary dipoles.
• Temporary dipoles can cause dipoles to form in neighboring molecules, called induced dipoles.
• More electrons in the molecule increase the chance of temporary dipoles forming, making London forces stronger and increasing the boiling point.
• The increasing boiling points of halogens down group 7 and alkanes is due to an increasing number of electrons.
• Long straight chain alkanes have a larger surface area of contact than branched alkanes, and so have stronger London forces.
• Permanent dipole-dipole forces occur between polar molecules and are stronger than London forces.
• Polar molecules are asymmetrical and have a bond with a significant difference in electronegativity.
• Hydrogen bonding occurs in compounds with a hydrogen atom attached to nitrogen, oxygen, or fluorine.
• These atoms must have an available lone pair of electrons.
• Hydrogen bonds have large electronegativity differences between H and O, N, or F. Showing lone electrons on the contributing atoms is key.
• Hydrogen bonding occurs in addition to London forces.
• Bond angles around the H atom equal 180° because of two pairs of electrons around the H atom repelling to a position of minimum repulsion.
• Alcohols, carboxylic acids, proteins, and amides can form hydrogen bonds.
• Water can form two hydrogen bonds because the electronegative oxygen atom has two lone pairs, leading to a higher boiling point.
• in ice the molecules are held further apart than in liquid water thus its lower density.
• Hydrogen bonding is stronger than the other two types of intermolecular bonding.
• Anomalously high boiling points of H2O, NH3 and HF are caused by hydrogen bonding in addition to London forces, which require more energy to break.

Solvents and Solubility

• Solubility is a balance of energy required to break bonds against the energy released making new bonds.
• When an ionic lattice dissolves in water, it involves breaking up the bonds in the lattice and forming new bonds between the ions and water molecules.
• Negative ions are attracted to the positive hydrogens on polar water molecules, and positive ions are attracted to the negative oxygen, which are described as the hydration of the ions.
• Higher charge density increases hydration enthalpy, attracting water molecules more strongly.
• Smaller alcohols are soluble in water because they can form hydrogen bonds with water, but longer hydrocarbon chains decrease solubility.
• Compounds that cannot form hydrogen bonds with water, such as halogenoalkanes or nonpolar substances like hexane, will be insoluble in water.
• Non polar solutes will dissolve in non-polar solvents, e.g. iodine (London forces) will dissolve in hexane (London forces).
• Propanone is a useful solvent because it has both polar and non polar characteristics.
• It can form London forces with some non-polar substances such as octane with its CH3 groups. Its polar C=O bond can also hydrogen bond with water.

Metallic Bonding

• Metals consist of giant lattices of metal ions in a sea of delocalised electrons.
• Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons.
• Factors affecting the strength of metallic bonding: number of protons, number of delocalised electrons, and size of ion.
• Metals have high melting points because of the strong electrostatic forces between positive ions and delocalised electrons.
• Mg has stronger metallic bonding than Na because Mg has more outer shell electrons so more electrons are released and the Mg ions are smaller than the Na. There is therefore a stronger electrostatic attraction between positive and delocalised electrons.
• Metals conduct electricity well because delocalised electrons can move through the structure.
• Metals are malleable because the positive ions in the lattice are identical: ions slide easily over one another and attractive forces remain the same.

Structure of Solids

• Structures can be ionic, metallic, molecular, or giant covalent (macromolecular).
• Ionic solids (giant ionic lattices), covalently bonded solids (diamond, graphite and silicon(IV) oxide), and solid metals (giant metallic lattices) are present in giant lattices.
• Ionic Sodium chloride exist as giant ionic lattice showing alternate Na (positive) and Cl (negative) ions.
• Giant covalent diamond has Tetrahedral arrangement of carbon atoms with 4 covalent bonds per atom.
• macromolecular graphite has Planar arrangement of carbon atoms in layers, with only 3 covalent bonds per carbon. The 4th outer electron per atom is delocalised between layers.
• Both giant structures have high melting points because of their many strong covalent bonds so a lot of energy is needed to break these.
• Metallic lattices are characterised with a giant metallic lattice showing close packing e.g. magnesium ions.
• Regular arrangements of I2 molecules held together by weak London forces.
• Molecular Ice has a central water molecule with two ordinary covalent bonds and two hydrogen bonds in a tetrahedral arrangement. The molecules in ice are held further apart than in liquid water and this explains the lower density of ice

Carbon Allotropes

• Macromolecular diamond has Tetrahedral arrangement of carbon atoms with 4 covalent bonds per atom. Diamond cannot conduct electricity well because all 4 outer electrons per carbon atoms are are involved in covalent bonds. They are localised and cannot move.
• Macromolecular Graphite has planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th outer electron per atom is delocalised between layers.
• Graphite can conduct electricity well between layers because one electron per carbon is free and delocalised, so electrons can move easily along layers.
• Does not conduct electricity from one layer to the next as electron transfer is too large.
• Both giant structures have high melting points because of their many strong covalent bonds so a lot of energy is needed to break these.
• Graphene has very high tensile strength because of the strong structure of many strong covalent bonds. This can conduct electricity well along the structure because one electron per carbon is free and delocalised, so electrons can move easily along the structure.
• Nanotubes have very high tensile strength because of the strong structure of many strong covalent bonds.
• Nanotubes can conduct electricity well along the structure because one electron per carbon is free and delocalised, so electrons can move easily along the structure.
• Nanotubes have potentially many uses, for delivery of drugs to cells.
• There are delocalized electrons in buckminsterfullerene.

Bonding and Structure Summary

• Ionic bonding involves electrostatic force of attraction between oppositely charged ions. The structure is a giant ionic lattice e.g. Sodium chloride and Magnesium oxide.
• Covalent bonding involves shared pair of electrons. Simple molecular structures involve weak intermolecular forces between molecules e.g. the molecules lodine, ice, carbon dioxide, water, and methane
• Further Covalent sharing results in giant molecular structures, where the structure as a whole is macromolecular. Examples incude Diamond, Graphite, Silicon dioxide, Silicon.
• Metallic structures involves forces of attraction between the metal positive ions and the delocalised electrons. and result in Giant metallic lattices structure. for example magnesium and sodium
• Use the terms molecules and intermolecular forces when talking about simple molecular substances.
• Ionic has high bp and mp because of strong electrostatic forces in giant lattice. Soluble in water, but the structure is poor conductivity as the ions fixed in the lattice.
• Simple molecular substances have low bp and p because of weak intermolecular forces between molecules and poor water solubility. Poor conductivity.
• Macromolecular substances have high bp and mp as high energy is needed to overcome the many strong covalent bonds. structure shows in-solubility in water. structure also shows poor electrical conductivity apart from silicon.
• Metallic substances have high bp and mp because of strong the strong electrostatic forces . structure is insoluble but a good electricity conductor.
• General physical description is that ionic = crystalline solids, molecular = gases and liquids, macromolecular = solids and metallic = shiny metal