Prof Gilbert LECTURE 21 CHEM 1211 10 28 10 Last time Chapter 7 Allowed combinations of Q N values Filling in orbitals Writing electron configurations This time Chapter 8 Chemical Bonding Periodic properties of the elements related to their atomic structures 1 Ionic bond the electrostatic energy Eel of attraction between ions of opposite charge in crystals of ionic compound such as NaCl How much energy depends on the product of the charges on the ions Q and Q divided by the distance d between them according to Coulomb s law Eel ke Q Q d where d is the sum of the radii of the ions for Na and Cl d equals 102 181 283 pm 2 Covalent bond Shared pairs of electrons between the atoms of non metals such as Br inside molecules and molecular ions Why do they form Two explanations 1 sharing lets atoms complete their octets that is acquire the electron configurations of noble gas atoms 2 Mutual attraction between the nuclei of bonding atoms and the electrons of the other atom same electrostatic explanation that we used for ionic bonds 3 Metallic bonds The nuclei of metallic atoms attract valence electrons of surrounding atoms These attractions result in the formation of metallic bonds Electrons are not shared between pairs of atoms instead they are completely delocalized forming a sea of mobile electrons that move freely between all the atoms in the metal The Octet Rule Atoms of many of the representative elements Groups 1 2 13 17 gain stability by sharing pairs of electrons to fill their valence shell s and p orbitals For Z 5 atoms this means acquiring an octet of 8 valence elections Lewis symbols of the elements Valence shell s and p electrons represented by dots Number of unpaired dots atom s bonding capacity Capacity 1 2 3 4 3 2 1 0 Lewis structures show how pairs of valence electrons are shared in bonds or remain as lone pairs on the atoms in molecules or molecular ions To determine how many bonds shared pairs there are we can use the S N A rule As long as the octet rule is obeyed the number of electrons Shared among the atoms in a molecule equals the number Needed for noble gas electron configuration less the number Available Drawing Lewis Structures some in class examples NH3 CO C2H5OH ethanol HNO3 CO3 2 Polar bonds Sharing electrons does not necessarily mean equal sharing Different elements have different attraction for the shared pairs In other words different elements have different electronegativities Electronegativity is expressed using a relative scale that goes from 0 8 Cs to 4 0 F Most metals have values of 1 0 1 9 metalloids are 2 0 2 4 and non metals are 2 5 4 0 Note the periodic trends in electronegativities the overall height of the bars of the first 20 elements less the noble gases in this figure and how they map onto the trends in ionization energies the heights of the darker color portions different electronegativities such as H and Cl share a pair of electrons the pair spends more time closer to the more electro negative atom Cl giving that end of the bond a negative polarity and the other H end a positive polarity Here are some ways to represent unequal sharing and bond polarity When two atoms with Ions may be considered extreme cases of unequally shared electrons Bonds are considered ionic rather than compounds form when the difference in electronegativity between the two atoms is 2 0 NOTE the difference between a polar covalent bond and an ionic bond is a matter of the degree of inequality in electron sharing and not clearly defined Resonance Sometimes a single Lewis structure does not adequately describe the bonding in a molecules of molecular ion We have already seen some examples HNO3 NO3 CO3 2 Here is one more ozone O3 has two identical bonds but its Lewis Structure requires one single and one double bond To rationalize the difference we assume that the single double bond structures resonate Resonance can happen when bonding and lone pairs are free to move without breaking any single bonds Choosing Between Lewis Structures Formal Charges You may be able to draw more than one Lewis structure for a molecule in which the octet rule is obeyed by all the atoms To determine which structure is best calculate the formal charge on each atom and determine which structure minimizes these charges that one is preferred How to calculate formula charges Formal charge of valence electrons on the free atom of valence electrons assigned to the atom in the structure In class example the molecular structure of N2O How do we choose between NOTE When the of bonds to an atom in a Lewis structure matches the atom s bonding capacity its formal charge FC is 0 If there is one fewer bonds then FC 1 if one more bond then FC 1 Execptions to the Octet Rule 1 Odd number of electrons Examples NO and NO2 Question 1 on which atom does the odd electron go 2 Which is the preferred resonance structure for NO2 2 Less than an octet Compounds of Be and B e g BeCl2 and BF3 3 More than an octet Many compounds and polyatomic ions with atoms from the 3rd period onwards can accommodate more than an octet because the d orbitals are low enough in energy to participate in bonding Consider SF6 SO4 2 PO4 3