1. 1 ATOMIC STABILITY Ion and Molecule Formation

2. 2 Electronegativity  In a covalent bond, we have seen that electron pairs are shared between two nonmetals  Rarely are these electrons shared equally as one of the atoms has a stronger “desire” to have those electrons  How can we measure which atoms wants the electrons more?  Electronegativity (EN)  It is a measure of an atoms ability to attract a pair of electrons in a molecule

3. 3  F is the most electronegative element and is given a value of 4.0 and all elements E.N. values are in comparison to this  Left to right across a period =  in EN  Down a group  in EN or stays about the same Electronegativity

4.  The higher the EN value, the more the atom will attract shared electrons to it  Depending upon how great the difference in electronegativity is between the atoms the bond can have highly positive and negative regions  This is called a polar bond 4

5. 5  The only bond that is purely 100% covalent where the electrons are equally shared is one in which the  EN = 0  This only occurs when the electrons are shared by identical atoms, like H 2, or any of the diatomic molecules  Every other bond will have a % ionic character and a % covalent character based upon the  E N Electronegativity

6.  A bond is considered to be non-polar covalent if the  EN is 0 – 0.3  A bond is considered to be polar-covalent if the  EN is 0.3 – 1.7  Any bond with  EN > than 1.7 is ionic in character 6

7. Electronegativity difference difference 0 0.3 > 1.7 Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases Nonpolar NonpolarCovalent 95% Covalent 5% Ionic 50% Covalent 50% Ionic 100% Covalent 1.7

8. Electronegativities of the Elements

9.  Using the chart of Electronegativities, determine the type or bond formed between the following pairs of atoms:  C and O  Fe and O  N and Br  C and H  Na and F Electronegativity

10. 10  In a molecule of H 2 O, a pair of electrons are shared between each O and H  The EN of O = 3.5 and H = 2.1   EN = 1.4 – therefore is a polar covalent bond  This means that O attracts the electrons towards it and so will become slightly negative while the electrons move away from each H atom and they become a bit positive  This means the electrons are not shared evenly and that one area is slightly positive, the other negative.  This is called a polar molecule  Indicated using small delta (δ). Electronegativity

11. Dipole Moments  A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),  or has a dipole moment.  Center of charge doesn’t have to be on an atom.  Will line up in the presence of an electric field.

12. How It is drawn H - F ++ --

13. ++ --

14. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

15. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- + -

16. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- - +

17. Which Molecules Have Dipole Moments?  Any two atom molecule with a polar bond. Eg. NaCl or FBr  With three or more atoms there are two considerations. 1. There must be a polar bond. 2. Geometry can’t cancel it out (more about geometry later) Eg. CH 4 CO 2 SO 2

18. 18 Ionic vs. Molecular Compounds  There are two types of forces involved in chemistry  Intermolecular forces are those between molecules and are responsible for holding these molecules together (inter = between)  Intramolecular forces are those between atoms inside the actual molecule and are responsible for holding the molecule together (intra = within)  These two forces explain many of the properties of ionic and covalent compounds

19. 19  Ionic compounds are formed of positive and negative ions and these forces are very strong  Each ion is held in place by at least 6 other ions and so both the inter and the intra molecular forces are strong  Covalent compounds have strong intramolecular forces holding the atoms together to form a molecule, but rather weak intermolecular forces holding the adjacent molecules together  Because the intermolecular forces are weak, covalent compounds have low boiling and melting points (little energy is needed to move molecules apart from a solid to liquid to gas)  Many are gases at room temp Ionic vs. Molecular Compounds

20. 20  Ionic compounds tend to have high melting points and boiling points as much energy is required to pull the ions apart  Most are solids at room temperature  Ionic compounds are also hard and brittle  This is because that while it is difficult to break the ions apart (hard) if the ions get moved slightly the ions line up to repel one another and force the solid to “break”  It is dependent upon how the crystal is arranged  Carbon can be arranged to form graphite (brittle) or diamonds (hard) and each had a different crystal arrangement using the same atoms Ionic vs. Molecular Compounds

21. 21  Solid ionic compounds do not conduct electricity as the ions are held tightly, but when in the liquid state (called molten) the ions are free to move and so can conduct electricity  Ionic compounds dissolve easily in water as water is a polar molecule and water molecules surround the ions and pull them apart into the solution (process called solvation) Ionic vs. Molecular Compounds

22. Metallic Bonding  How are metal atoms held in place?  Most metals have 1, 2 or 3 valence electrons  The metal atoms are relatively close to each other and their valence energy levels overlap  This allows the valence electrons to move freely from one metal atom to those it overlaps with  These electrons are not bonded to one particular metal atom and are called delocalized electrons  This is often referred to as the “Electron Sea Model” of metallic bonding 22

23. Ionic Bond, A Sea of Electrons 23 Metallic Bonding

24.  Because each metal atom is sort of positive and is surrounded by this moving “sea” of e -, the atoms are held in place very strongly  The bond strength is variable and depends upon the number of electrons each atom allows to roam in the “sea” and thus the sea size is the determining factor  Fe & Ti have large seas and so are strong metals  The sea of e - also explains the conductivity of metals as e - are free to move through the metal  The ductility and malleability of metals is explained by the fact that metal atoms are separated easily and in all directions 24 Metallic Bonding

25. Sea of Electrons  Metals conduct electricity.  Electrons are free to move through the solid. 25

26. Metals are Malleable  Hammered into shape (bend).  Also ductile - drawn into wires.  Both malleability and ductility explained in terms of the mobility of the valence electrons 26

27. Malleable ++++ ++++ ++++  Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. Force 27

28. Ionic solids are brittle +-+- + - +- +-+- + - +- Force 28

29. Ionic solids are brittle + - + - + - +- +-+- + - +-  Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. Force 29

30. Hydrogen Bonding  This is a type of bonding involving hydrogen and either F, O or N  When hydrogen bonds with either of these elements there is a large ΔEN  This results in a very polar molecule with large dipoles  This produces relatively high inter molecular forces to adjacent molecules they are held together “tightly”  This accounts for the relatively high boiling and melting point of H 2 O compared to other covalent compounds (table of bp’s on page 190) 30

31. Hydrogen Bond 11.2 The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. 31