1. Section 6.5 – Molecular Compounds Binary molecular compounds are composed of two nonmetallic atoms. Because atoms can combine in different ratios (for example CO and CO 2 ) we use prefixes to help distinguish between compounds. CO is carbon monoxide CO 2 is carbon dioxide CCl 4 is carbon tetrachloride Note the –ide ending (similar to how an anion works, but these aren’t ionic compounds) 1 PrefixNumber mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 deca-10

2. Section 6.5 – Molecular Compound Naming To convert a name to a formula, write the correct symbols for the two elements, then add appropriate subscripts. If there is just one of the first atom, you don’t need to write the mono-, it is assumed. But for the second atom, if there is one, use mono- Ex: tetraiodine nonoxide is I 4 O 9 sulfur trioxide SO 3 phosphorus pentafluoridePF 5 Ex: N 2 O dinitrogen monoxide PCl 3 phosphorus trichloride SF 6 sulfur hexafluoride H 2 O dihydrogen monoxide 2

3. Molecular Bonding - Acids Here is a list of some of the most common acids which have covalent bonds and their names, which don’t always follow the standard naming convention. HCl Hydrochloric acid H 2 SO 4 Sulfuric acid HNO 3 Nitric acid CH 3 COOH Acetic acid (also written HC 2 H 3 O 2 ) H 3 PO 4 Phosphoric acid H 2 CO 3 Carbonic acid These are the most common ones and good to memorize 3

4. Single Covalent Bonds 4 Hydrogen is the simplest model of a covalent bond Each Hydrogen has one electron and they share them to form a single covalent bond. The single covalent bond can be represented by the pair of electrons or as a dash as shown below H:H or H-H

5. Conventions for naming 5 The chemical formulas of ionic compounds describe _________________________ (Example: NaCl is a formula unit) The chemical formulas of covalent compounds describe ________________________. (Example H 2 O is a molecule) Ionic compounds do not have molecular formulas because they are not composed of molecules.

6. Covalent Molecules 6 Combinations of atoms of the nonmetallic elements in groups 4A, 5A, 6A and 7A of the periodic table are likely to form ___________________. Chemist Gilbert Lewis summarized this tendency in his formulation of the _____________________________________ _____________________________________ Sharing of electrons occurs if the atoms involved acquire the electron configuration of _______________________.

7. Covalent Bonding – Diatomic Gas Fluorine 7

8. In a water molecule, two hydrogen atoms form one single covalent bond each with one oxygen atom. Note how the O atom ends up with eight electrons around it. Covalent molecules will form if each atom will end up with 8 electrons around it (except H). Each dot is ______________. Each line is ______________. 8

9. 9 Three Single Bonds Four Single Bonds

10. The HONC Rule  HH  Hydrogen (and Halogens) form _____________________  O  Oxygen (and sulfur) form __________________________  One double bond, or two single bonds  N  Nitrogen (and phosphorus) form ____________________  One triple bond, or three single bonds, or one double bond and a single bond  C  Carbon (and silicon) form _________________________.  Two double bonds, or four single bonds, or a triple and a single, or a double and two singles 10

11. Review: The Octet Rule and Covalent Compounds 11  Covalent compounds tend to form so that each atom, by _____________ electrons, has an ____________ of electrons in its highest occupied energy level.  Covalent compounds involve atoms of _____________________________.  The term “__________________________” is used exclusively for covalent bonding

12. Coordinate Covalent Bonds 12 A covalent bond in which ______________________ __________________________________________ __________________________________________. This is signified by showing coordinate covalent bonds as ___________ that point from the atom donating the pair of electrons to the atom receiving the bond. Many polyatomic cations and anions contain both covalent and coordinate bonds. NH 4 + is an example.

13. Resonance  Occurs when more than one valid Lewis structure can be written for a particular molecule, such as ozone, below. 13  These are resonance structures. The actual structure is an average or a blend of the resonance structures.

14. Resonance in Benzene, C 6 H 6 14 Each of these junctions represents where a Carbon is

15. Section 16.3: Polar bonds and molecules Covalent bonds involve sharing electrons between ____________________________ Sometimes the sharing is equal and the electron resides halfway in between the atoms, as in a diatomic gas like N 2, Cl 2, etc. This is called a ___________________________________________ If the bonding electrons are shared unequally, this is called a ______________________________________________ 15

16. Polar Bonds The greater the electronegativity value, _________________ _________________________________________________. A high electronegativity atom is not “stealing” electrons as in the ionic case, but it is moving them in its direction. Consider HCl. Hydrogen has an electronegativity of 2.1 and Chlorine has an electronegativity of 3.0. These values are quite different, so the covalent bond in HCl is polar. The shared electrons are pulled in the direction of Cl, because it is more electronegative. This can be represented as follows: 16

17. Bond Polarity Electronegativity Differences and Bond Types Electronegativity Difference Range Most probable type of bond Example 0.0 – 0.4Nonpolar covalentH – H (2.1-2.1=0.0) 0.4 – 1.0Moderately polar covalent d+ d- H – Cl (3.0-2.1=0.9) 1.0 – 2.0Very polar covalentd+ d- H – F (4.0-2.1=1.9) ≥ 2.0IonicNa + Cl - (3.0-0.9=2.1) 17

18. Polar Molecules Some molecules have dipoles (aka dipolar molecules), but their polarities line up in such a way that they ________________________. Carbon dioxide is one such example 18

19. Intermolecular Attractions- van der Waals forces The weakest intermolecular attractions are van der Waals forces. These consist of two possible types, London dispersion forces and dipole interactions. London dispersion forces, (weakest of all intermolecular interactions) are caused by the motion of electrons. The strength of dispersion forces increases as the number of electrons increases. For halogens, which have more e- in their outer shell, the major attraction between them is dispersion forces. These forces are weaker for F and Cl (gases at STP). They are stronger for Bromine, a liquid at STP, and even stronger for Iodine, a solid at STP. 19

20. Dipole interaction forces The second type of van der Waals force is the dipole interaction, when polar molecules are attracted to one another. The positive region of one molecule is attracted to the negative region of another. 20 HCl molecules

21. Hydrogen Bonds A hydrogen bond is an attractive force where a hydrogen which is covalently bonded to a very electronegative atom (meaning the H has a slight δ + charge on it) is also weakly bonded to an unshared electron pair of another atom (pair has δ – charge to it). This happens because when H bonds to O, F or N, the very polar bond leaves the H very electron deficient, with essentially an exposed nucleus (a proton) with no electrons. The H nucleus is then attracted to a negatively charged unshared electron pair on another atom. The resulting hydrogen bond is only about 5% of the strength of a regular covalent bond, but it is still the strongest of the intermolecular forces. – This is what causes water to be a liquid at room temperature 21

22. Hydrogen bonds in water 22 Hydrogen has valence e- that are not shielded from the nucleus by another layer of electrons. Water has this type of interaction because the hydrogens have a slightly + charge and the oxygen has a slightly – charge. This relatively strong interaction is called a hydrogen bond.

23. Intermolecular Forces Summary Weakest – Middle – Strongest – But all three are still much weaker than a 23

24. Intermolecular Attractions and Molecular Properties The physical properties of a compound depend on the type of bonding it has – ____________________________________. Here are some comparisons of physical properties 24 Characteristics of Ionic and Covalent Compounds CharacteristicIonic CompoundCovalent Compound Representative unit Bond formationTransfer of one or more electrons between atoms Sharing of electron pairs between atoms Type of elements Physical stateSolidSolid, liquid and gas Melting pointHigh (usually > 300 C)Low (usually < 300 C) Solubility in waterUsually highHigh to low Electrical conductivity of aqueous solution Good conductorPoor conductor or nonconducting