Topic 4: Bonding Adapted from Mr Field
4.1 Ionic Bonding and Structure
Ionic bonding An ionic bond is: The electrostatic attraction between two oppositely charged ions sodium fluoride lithium oxide F - Na + Li + O 2- Li + Ionic bonds typically form between a metal and a non-metal Ionically bonded compounds are often referred to as salts
Ionic Character You can predict ionic character by two ways: Position on the Periodic Table Electronegativity differences Position on the Periodic Table Metals on the left side of the zigzag line tend to lose electrons. Non-metals on the right side of the zigzag line tend to gain electrons. The highest tendency to form ionic compounds will be between elements on the bottom left with elements on the top right. Electronegativity Electronegativity values are given in Table 8 of the IB Data Booklet. Differences greater than 1.8 are considered to be ionic (more on this later)
Formation of simple ions Positive ions (cations) Positive ions are formed when metals lose their outer shell electrons Group 1: Li Li + + e - Group 2: Ca Ca 2+ + 2e - Group 3: Al Al 3+ + 3e - Transition metals form multiple different ions Fe Fe 2+ + 2e - Fe Fe 3+ + 3e - Negative ions (anions) Negative ions are formed when non-metals gain enough electrons to complete their outer shells Group 5: N + 3e - N 3- Group 6: O + 2e - O 2- Group 7: F + e- F -
Polyatomic ions Many ions are made of multiple atoms with an overall negative charge The negative ones are mostly acids that have lost their hydrogens You need to know about: Sulphate, SO 2-4 Phosphate, PO 3-4 Nitrate, NO - 3 Carbonate, CO 2-3 Hydrogen carbonate, HCO - 3 Hydroxide, OH - Cyanide, CN - Ammonium, NH + 4
The names of ionic compounds The cation gives the first part of the name Normally a metal except in the case of ammonium (NH 4+ ) In the case of transition metals, Roman numerals tell you the charge on the metal ion The anion gives the second part of the name Simple ions: -ide e.g. chloride, fluoride, nitride etc Complex ions: just their name: sulphate, phosphate etc Note: the -ate ending usually refers to polyatomic ions containing oxygen, which provides the negativity more on this in the redox unit Examples: CaF 2 : calcium fluoride Fe 3 (PO 4 ) 2 : iron (II) phosphate
Lattice Structure Under normal circumstances, ionic compounds are usually solids with a lattice structure The ions in ionic compounds surround themselves with ions of the opposite charge. Therefore ionic compounds have three-dimensional crystalline structures known as ionic lattices. The strength of force between the ions is known as lattice enthalpy and depends upon the size of the ion and the charge on the ion. The smaller the ion and the greater the charge, the higher the lattice enthalpy. The formula unit is the smallest particle of an ionic compound.
The formula of ionic compounds Ionic compounds are always neutral, so the charges must balance Example 1: calcium reacting with fluorine: Calcium forms Ca 2+, fluorine forms F - The formula is CaF 2 so two F - charges cancel the one Ca 2+ Example 2: iron (II) reacting with phosphate Iron (II) is the Fe 2+ ion, phosphate is PO 4 3- The formula is Fe 3 (PO 4 ) 2 The 6 + charges from iron (2 + x 3) balance the 6 - charges (3 - x 2) from phosphate Look for the lowest common multiple Ionic compounds do not form molecules so these are always empirical formulae
PHYSICAL PROPERTIES Remember physical properties can be observed without chemically altering the substance. MELTING AND BOILING POINTS Ionic compounds have high melting and boiling points because the electrostatic attraction between the ions is very strong and requires large amounts of energy to break the bond. The higher the charge and smaller the ion, the greater the melting point. VOLATILITY: the tendency of a substance to vaporize. Ionic compounds have very low volatility.
PHYSICAL PROPERTIES conti. SOLUBILITY: the ease with which the solid dissolves in a liquid to become a solution. Solubility trends are based on the similarity of the chemical nature of the solute and the solvent. Polar compounds dissolve in polar solvents. Nonpolar compounds dissolve in nonpolar solvents Like dissolves like
PHYSICAL PROPERTIES conti. ELECTRICAL CONDUCTIVITY: is the ability to conduct electricity. Freely moving ions must be present to conduct electricity. Molten or dissolved ionic compounds conduct electricity. Solid ionic compounds do not because the ions are locked into place and are not free to move about. BRITTLENESS: means the crystal will shatter when force is applied. Ionic compounds tend to be brittle because ions of like charge can be next to each other in the lattice structure and the repulsive charges cause the structure to split easily.
Key Points Ionic bonds are the attraction between two oppositely charged ions Ionic bonds form between metals and non metals Metals lose their outer shell Non-metals complete their outer shell The number of each ion in the formula is determined by the lowest common multiple of their charges
4.2 Covalent Bonding
Lesson 2: Covalent Bonding Objectives: Refresh knowledge and understanding of covalent bonding Learn how to draw Lewis structures Identify examples of coordinate covalent bonding Identify instances of expanded octets
Covalent bonding A covalent bond is the attraction of two atoms to a shared pair of electrons water Each O has two single bonds carbon dioxide each C has two double bonds H O H O C O Atoms aim for complete outer-shells, and each covalent bonds gives them one electron Atoms form as many bonds as they have gaps in their outer-shells Covalent bonds typically form between two non-metals
ELECTRON SHARING When atoms of 2 non-metals react together, each is seeking to gain electrons in order to achieve the stable electron configuration of a noble gas. This tendency to form a stable arrangement of 8 electrons in the outer shell is referred to as the octet rule. The shared pair of electrons is concentrated in the region between the 2 positively charged nuclei. The electrostatic attraction between the 2 nuclei and the electrons constitutes the covalent bond.
How many bonds? Atoms (usually) form bonds according to the octet rule This means they try to get a full outer shell of 8 electrons (except hydrogen which is full at 2) Atoms form as many bonds as they have gaps in their outer shells, with each bond gaining them one electron: Group 7: 7 electrons, 1 gap 1 bond Group 6: 6 electrons, 2 gaps 2 bonds Group 5: 5 electrons, 3 gaps 3 bonds Group 0/8: 8 electrons, 0 gaps 0 bonds
Multiple Bonds Covalent bonds can be: Single: one shared electron pair, X-X Double: two shared electron pairs, X=X Triple: three shared electron pairs, XX Triple bonds are stronger than double bonds which are stronger than single bonds. The strength of the bond is a measure of how much energy is required to break the bond. Triple bonds are shorter than double bonds which are shorter than single bonds. The number of shared electrons is greater in multiple bonds causing the electrostatic attraction to be stronger; therefore, causing the bonds to be shorter in length.
Covalent Character You can predict covalent character by two ways: Position on the Periodic Table Covalent compounds tend to form between 2 non-metals. The closer together two elements are, the more covalent. Electronegativity Electronegativity values are given in Table 8 of the IB Data Booklet. Differences less than 1.8 are considered to be covalent.
Polarity A bond that is unsymmetrical with respect to electron distribution is said to be polar. The term dipole means the bond has an area of positive charge and an area of negative charge. (delta symbol)
Polarity The Pauling scale predicts polarity by using electronegativity differences. Values of 0.0-0.4 are considered non-polar covalent. Values between 0.4 and 1.8 are considered polar covalent. Values greater than and equal to 1.8 are considered ionic. Basically, the greater the electronegativity difference, the more polar the bond.
4.3 Covalent Structures
Covalent Structures Molecular As in water and methane Giant lattice As in silicon dioxide More on these later in the unit
Lewis structures Show the position of outer-shell electrons in a covalent compound Various types: all show the same thing, any is fine dots and crosses crosses only dots only lines Blue Circles: These are the bonding pairs of electrons the ones involved in the bonds. Red Circles: These are non-bonding or lone pairs of electrons. They are very important, but students often forget about them!
Rules for Lewis Structures Add up the total number of valence electrons in the molecule. Draw the skeletal structure. Use a line between each element to symbolize an electron pair. Distribute the remaining electrons around the elements in pairs to form octets. (Hydrogen can only ever have 2 electrons.) If you do not have enough to form octets, make double or triple bonds. Ions must have square brackets around them with the charge notated in the top right hand corner. To be a correct Lewis structure, ALL electrons must be shown.
Reduced Octet Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.
The expanded octet In this example, the Lewis structure of PCl5 shows it with 10 electrons in the outer shell This is because phosphorus can make use of its empty d-orbitals (the 3d ones) This is called an expanded octet Period 2 elements can t do this as they have no d-orbitals
Resonance Resonance is a concept used to describe the structures when there are multiple ways to depict the same molecule. If you can put a double bond in more than one position, you will be expected to draw the resonance structures. The electrons are actually delocalized in the areas of the double bonds and are spread out equally among all bonding positions. Bond strength and length are in between that of single and double bonds.
Resonance structures allow us to depict all the possible positions of the double bonds. The true structure, however, is an intermediate form known as a resonance hybrid. Double arrows are placed between all resonance structures. Ref: myweb.astate.edu
The Coordinate Covalent/Dative-bond Sometimes an atom will contribute both of the electrons in a covalent bond, this is called a dative (covalent) bond E.g. In this example, the lone pair from a water molecule has formed a coordinate covalent bond to a hydrogen ion (H + ) You can show dative bonds with an arrow to say where the electrons came from but do not have to
More examples of coordinate covalent bonds
Octet Rule Exception Hydrogen will never have more than 2 electrons. Be and B have less than 8 electrons. Some elements like S and P can have expanded octets which hold more than 8 electrons. Coordinate covalent bonds are formed when both electrons originate from the same atom. An arrow is used to denote the direction in a coordinate covalent bond showing the atom from which both electrons originated. This is common in double and triple bonds.
Key Points Atoms (generally) form covalent bonds according to the octet rule. Each covalent bond gives an atom one extra electron Period 3 (and above) elements can break the octet rule by using empty d-orbitals and might have 12 or more electrons in their outer shell
Molecular Shapes Shapes of species are determined by the repulsion of electron pairs according to the VSEPR theory Objective: Be able to use the VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains.
VSEPR a brief introduction Valence Shell Electron Pair Repulsion (aka vesper ) Pairs of electrons around an atom repel each other this determines a molecule s shape Pairs of electrons are known as charge centres and include both: The electrons in a covalent bond a double/triple bond only counts as one charge centre! Lone pairs / non-bonding pairs Example: ethyne The carbon has two charge centres (the C-H bond and the CC bond) They push as far away from each other as possible making a 180 o bond angle
Key Points Negative charge centres repel each other, this determines the shape of a molecule Standard shapes: Two charge centres: linear Three charge centres: trigonal planar Four charge centres: tetrahedral Five charge centres: trigonal bipyramidal Six charge centres: octahedral
Drawing shapes in three dimensions: Draw as many atoms as you can in the same plane flat on the paper Use solid wedges to show atoms coming out of the plane of the paper towards you Use dashed wedges to show atoms going back into the plane of the paper away from you
Polarity The polarity of a molecule depends upon: The polar bonds it contains. The shape of the molecule. If the bonds are equally polar and arranged symmetrically, then they cancel each other out and are non-polar. If the molecule contains bonds of different polarities or the bonds are not arranged symmetrically, then the molecule will be polar. You can usually tell by the shape and lone pairs of electrons if the molecule is polar or not.
Electron Domains Double and triple bonded electron pairs are orientated together and behave as a single unit known as an electron domain. Lone pairs also count as electron domains so the 12 VSEPR shapes are narrowed down to 5 basic shapes. For the electron domain shapes, you will need to know bond angles.
LINEAR A linear molecule has two electron domains. The angle is 180 degrees Non polar
TRIGONAL PLANAR (3 electron domains) A trigonal planar molecule has 3 electron domains. It has angles of 120 degrees. Non polar The bent molecule can also have 3 electron domains (2 shared and 1 unshared) Polar
TETRAHEDRAL (4 electron domains) A tetrahedral molecule has four electron domains. It has angles of 109.5 degrees. Non polar Trigonal pyramidal (3 shared and 1 unshared) Polar Bent (2 shared and 2 unshared pairs) Polar
Allotropes of Carbon Some covalent structures are crystalline in nature like ionic lattices; however, they are linked together with covalent bonds. The crystal is a single molecule with a regular repeating pattern of covalent bonds. It is referred to as a giant molecular structure. Allotropes are different forms of an element in the same physical state. Different bonding within the structures gives rise to different physical properties. Carbon has four allotropes.
GRAPHITE Graphite Each C atom is covalently bonded to 3 others forming hexagons in parallel layers with bond angles of 120 degrees. The layers are held together by van der Waal s forces so they can slide over each other. Density is 2.26 g/cm 3 Contains one non-bonded, delocalized electron per atom so graphite conducts electricity due to the movement of these electrons. Not a good heat conductor Very high melting point, most stable allotrope Non lustrous grey solid Used as a lubricant and in pencils
DIAMOND Diamond Each C atom is covalently bonded to 4 others tetrahedrally in a regular repeating pattern with bond angles of 109.5 degrees. It is the hardest known natural substance. Density is 3.51 g/cm 3 All electrons are bonded so it does not conduct electricity. Does conduct heat better than metals. Very high melting point, brittle Lustrous crystal Used in jewelry and tools
FULLERENE, C 60 Fullerene Each C atom is covalently bonded in a sphere of 60 carbon atoms consisting of 12 pentagons and 20 hexagons. The structure is a closed spherical cage in which each carbon atom is bonded to 3 others. Density is 1.726 g/cm 3 It easily accepts electrons to form negative ions so it is a semiconductor at normal temp and pressure due to some electron mobility. Very low heat conductivity Low melting point Yellow crystalline solid soluble in benzene Related forms are used to make nanotubes for the electronics industry, catalysts and lubricants.
GRAPHENE Graphene Each C atom is covalently bonded to three other carbons forming hexagons with bond angles of 120. The structure is a two dimensional single layer described as a honeycomb or chicken wire structure. Density is 1.5 g/cm 3 Contains one non-bonded, delocalized electron per atom so graphene conducts electricity due to the movement of these electrons. Excellent heat conductor better than diamonds Very high melting point, thinnest material to exist, stronger than steel Almost completely transparent Used in touch screens, high performance electronics, etc
Si and SiO 2 Silicon is a group 4 element with 4 valence electrons. In its elemental state, each silicon atom is covalently bonded to four others in a tetrahedral arrangement. This results in a giant lattice structure like a diamond. SiO 2 commonly known as silica or quartz, forms a giant covalent structure. It is also a tetrahedral structure, but the bonds are between Si and O. Each Si atom is covalently bonded to 4 oxygen atoms and each O to two Si atoms. The formula SiO 2 is the ratio of atoms within the giant molecule. The structure is strong, insoluble in water, does not conduct electricity or heat and has a high melting point all properties of glass and sand which are different forms of silica.
Key Points Carbon (diamond, graphite and fullerenes), silicon and silicon dioxide exhibit giant covalent (macromolecular) structures For example diamond: each carbon is bonded to exactly four others and so on Ionic compounds form giant ionic lattices NaCl: every Na + ion surrounded by 6 Cl - ions, every Cl - ion surrounded by 6 Na + ions
4.4 Intermolecular Forces
Intermolecular Forces Objectives: Learn to identify and explain the three types of intermolecular forces: Van der Waals Permanent dipole-dipole Hydrogen bonds Understand and explain the effects of the above on melting/boiling points
Inter vs. Intra Intramolecular forces are forces within the molecule such as covalent, ionic and metallic bonding. Intermolecular forces are forces between molecules. The strength of intermolecular forces determines the volatility of a substance. The stronger the forces, the higher the melting and boiling points.
Intermolecular Forces The attractive forces between molecules It is these that are partially broken during melting, and fully broken during boiling Note: when molecular compounds melt/boil, the bonds in the molecule do not break, it is just the attractive forces between the molecules that break
Van der Waals Forces aka Temporary dipole induced dipole forces Random electron movements create a small, temporary dipole This induces a similar dipole in a neighbouring molecule This creates a small attraction between them These are weak and exist only for the tiniest fraction of a second Van der Waals forces are present in all molecules Van der Waals forces: Increase with molecular mass
Dipole-dipole forces aka Permanent dipole forces Different atoms have different electronegativities, which means there will be variations in the electron charge density in different parts of a molecule If a molecule is not symmetrical, the variation produces a dipole where a molecule as a positive and a negative end The end with high charge density is - The end with low charge density is + Oppositely charged dipoles attract each other. This is a relatively strong attractive force If a molecule is symmetrical, variations in electron charge density cancel each other out and the molecule is non-polar - + - + -
Hydrogen bonds aka H-bonds The strongest type of intermolecular force They occur between a nitrogen, oxygen or fluorine and a hydrogen that is bonded to a nitrogen, oxygen or fluorine N, O and F are the three most electronegative elements, and all have lone-pairs when bonded When H is bonded to N, O or F, the electrons in the bonded are strongly attracted to the N/O/F, leaving the H very positive The lone pair on the N/O/F is strongly attracted to the positive hydrogen
Effects of intermolecular forces Intermolecular forces play an important role in the properties of compounds including: Melting/boiling point: Stronger intermolecular forces higher m.p./b.p. Volatility: Stronger intermolecular forces lower volatility Solubility: like dissolves in like Polar solutes dissolve best in polar solvents Non-polar solutes dissolve best in non-polar solvents
Summary Three types of intermolecular force, from strongest to weakest: Hydrogen bonds Between N/O/F and H attached to N/O/F Dipole-dipole Between permanent dipoles on asymmetric molecules Van der Waals Between instantaneous dipoles formed on any molecule/atom
All molecules will have some type of van der Waal s force. Non-polar molecules have only London dispersion forces. Polar molecules have dipole-dipole forces and London dispersion forces. Hydrogen bonding exists when the positive hydrogen bonds with lone pairs of electrons on nitrogen, oxygen and fluorine.
4.5 Metallic Bonding
METALLIC BONDING Metallic bonding occurs when the delocalized electrons in metal atoms are attracted to the lattice of positive metal ions. Delocalized means that the electrons do not belong to any one metal nucleus but can spread themselves throughout the metal structure. Because the outer shell electrons in metals can easily be lost, positive ions are easily formed. When the metal atoms are together, the positive ions form a lattice where the electrons can freely move around.
The strength of the metallic bond depends upon three things: 1. The number of delocalized electrons 2. The size of the cation 3. The charge of the cation The number of delocalized electrons are the number of electrons in the outer shell. Sodium has one delocalized electron Magnesium has two delocalized electrons The greater the number of delocalized electrons and the smaller the cation, the greater the strength of the metallic bond.
The higher the charge, the greater the metallic strength. The strength of metallic bonding tends to decrease down a group as cations increase in size. Transition metals have very strong metallic bonds due to the large number of electrons that can become delocalized from both the 3d and 4s sub-levels.
PROPERTIES OF METALS Metals are good conductors of electricity and heat (thermal conductivity) because the delocalized electrons are highly mobile and move through the metal structure. They are also good conductors of heat due to the same electron mobility. Metals are malleable which means they can be shaped under pressure. This is due to the fact that the movement of electrons is non-directional.
Metals can also be ductile which means to be drawn into threads. Metals have high melting points because lots of energy is required to break the strong metallic bonds. Metals are shiny and lustrous because the delocalized electrons in the metallic structure reflect light.
MELTING POINTS Melting points tend to decrease down a group due to the reduced attraction of the delocalized electrons and the positively charged cations. The stronger the metallic bond (due to size of cation and charge), the stronger the melting point.
ALLOYS Alloys are solid solutions. Remember that a solution is a homogeneous mixture of a solute with a solvent. In alloys, the solute and the solvent are both solids. Two metals are mixed together in the molten state. When they solidify, ions of the different metals are scattered throughout the lattice and are bound by the delocalized electrons.
PROPERTIES OF ALLOYS Alloys are possible because of the non-directional nature of the delocalized electrons and the fact that the lattice can accommodate ions of different sizes. Alloys have properties that are distinct from their component elements due to the different packing of the cations in the lattice. The alloy is often more chemically stable, stronger and resistant to corrosion.
COMMON ALLOYS Steel iron with carbon and other elements High tensile strength but tends to corrode Stainless steel iron with nickel or chromium High strength and corrosion resistant Brass copper and zinc Plumbing Bronze copper and tin Coins and tools Pewter tin, antimony and copper Decorative objects Sterling silver silver and copper Jewelry and art objects