Covalent Bonding PDF

MED-102 General Chemistry Covalent Bonding  LOBs covered Describe the nature of a covalent bond Discuss the concept of electronegativity in the formation of polar bonds Determine whether a molecule is polar or not  Covalent Bond Some atoms prefer not to lose...

MED-102 General Chemistry Covalent Bonding  LOBs covered Describe the nature of a covalent bond Discuss the concept of electronegativity in the formation of polar bonds Determine whether a molecule is polar or not  Covalent Bond Some atoms prefer not to lose or gain electrons They can fill their outer shell by sharing electrons with other atoms Example: Formation of H2  Covalent Bond Each H atom has an 1s1 electronic configuration Each H atom needs one electron to form a stable [He] noble gas configuration Two H atoms come close and share each other’s electrons  Covalent Bond – Revision Slide WATCH: https://www.youtube.com/watch?v=qc1TGwedSIc Note that the 2 shared electrons are exactly in-between the two positive nuclei. The positive nucleus is attracted to the negative electrons in the middle. Therefore, both nuclei cannot move away. The Coulombic attraction to the negative electrons keeps the nuclei bound to each other. This is exactly how a covalent bond works. Two shared electrons constitute a single covalent bond. They can be represented by two dots, or by a single dash.  Covalent Bonding – O2 Each oxygen atom needs 2 electrons to fill its outer shell Two oxygen atoms combine together by overlapping their valence shells and sharing a total of 4 electrons – double bond  Covalent bonding – N2 Each N atom needs 3 electrons to fill its outer shell Two N atoms join together by overlapping their valence shells and sharing a total of 6 electrons Triple bond  Covalent Bond – Water Molecule Oxygen needs 2 electrons for a full outer n = 2 shell Hydrogen needs 1 electron to fill its n = 1 shell 1 Oxygen atom combines with 2 Hydrogen atoms  Covalent Bond – Methane, CH4 Carbon needs 4 electrons for a full outer shell Hydrogen needs only 1 Therefore 4 hydrogens and 1 carbon join together to form four single covalent bonds  Covalent bonding – NH3 Each N atom needs 3 electrons to fill its outer shell Each H atom needs 1 electrons Therefore 3 H atoms join together with 1 N atom to form ammonia, NH3  Covalent bonding – CO2 Each C atom needs 4 electrons to fill its outer shell Each O atom needs 2 electrons Therefore 1 C atom joins with 2 O atoms like shown below  Summary for Revision In ionic bonding, electrons are transferred from one atom to another so that all atoms can achieve octets. This forms ions of opposite charges, and the attraction between the ions is called an ionic bond. Ionic bonds usually involve a metal and a nonmetal. Nonmetallic elements can join together in a different way. Instead of transferring electrons between them, they get close to each other such that their valence shells overlap. This allows for a sharing of electrons, and the net result is that both atoms achieve octets. The shared electrons in the middle prevent the positive nuclei from moving away. This effect gives rise to what is known as a covalent bond. A single covalent bond consists of two shared electrons. It is represented by two dots, or by a single dash. A double covalent bond consists of four shared electrons. It is represented by four dots or by two dashes. A triple covalent bond consists of six shared electrons. It is represented by six dots or by three dashes. After forming covalent bonds, some atoms have unshared electrons in their outer shells. These unshared electrons come in pairs, and are called lone pairs of electrons. They are shown as two dots on the atom.  Electronegativity The power of an atom in a molecule to attract the bonding electrons towards itself Electronegativity determines how shared electrons are distributed Electronegativities to remember H 2.1 C 2.5 N 3.0 O 3.5 F 4.0 Cl 3.0 Br 2.8 I 2.5 5-Minute Break  Electronegativity Two like atoms Two different atoms 3.0 3.0 2.1 3.0 Atoms pull at the shared Chlorine atom pulls at the shared electrons with the same strength electrons with a greater strength Non-polar bond Polar bond  Electronegativity – ionic character Plot of electronegativity difference and percent ionic character  Electronegativity – ionic character – Revision Slide Consider the electronegativities of two connected atoms. If they have the same electronegativities, then the difference in electronegativities is zero (x-axis). On the y-axis, this corresponds to zero % ionic character. This means that we have a pure nonpolar covalent bond involved. If the difference in electronegativity between the two connected atoms is 1.7, then we find the y-axis value on the purple curve to be 50%. This means that the bond is 50% ionic and 50% covalent. This curve indicates that all polar covalent bonds always have some degree of ionic character. The curve also indicates that all ionic bonds always have some degree of covalent character. They cannot be 100% ionic. WATCH: https://www.youtube.com/watch?v=Ur1IAHMB4Uw  Exercise on Percent Ionic Character Consider the HCl molecule. What is the percent ionic character in HCl? H = 2.1 Cl = 3.0  Exercise on Percent Ionic Character Consider the HCl molecule. What is the percent ionic character in HCl? H electronegativity = 2.1 Cl electronegativity = 3.0 Difference in electronegativity = 3.0 – 2.1 = 0.9 (see x-axis) Tracing the red lines in the diagram above, we see that this corresponds to approximately 15% ionic character in HCl. Therefore, 85% covalent character.  Electronegativity Three different situations  Polar Covalent Bonds Notation and symbols 2.5 4.0 + and - are called partial charges They arise because all polar covalent bonds have some degree of ionic character In the above example, the difference in electronegativity is 4.0 – 2.5 = 1.5, and the purple curve we discussed previously gives approximately 40% ionic character.  Polar and nonpolar molecules You can have polar bonds but the molecule may be nonpolar  Polar and nonpolar molecules – Revision Slide In CO2, we see that we have two separate C – O bonds. Since O has a higher electronegativity than C, each bond is polar, and the bond dipoles (red arrows) point towards the more electronegative O atoms. These two C – O bond dipoles are like vectors, having a magnitude and direction. Since the two C – O bonds are identical, the magnitudes of the vectors are equal. But, the two vectors are pointing in exactly opposite directions. When we add the two vectors together, they cancel out completely. This means that the CO2 molecule does not have an overall polarity, it is nonpolar. In H2O, we have two O – H bonds. Since O is more electronegative than H, the bond dipoles are pointing towards the O atoms. The two O – H bond dipoles are both pointing slightly upwards. This means that when we add them together, they will not cancel out. Thus, the H2O molecule has a finite (nonzero) net dipole. It is a polar molecule.  Dipole moment A polar molecule has a dipole moment The higher the value of the dipole moment, the more polar the molecule A measure of polarity Important for deciding what solvent to use to dissolve a substance Polar substances dissolve in polar solvents Non-polar substances dissolve in non-polar solvents This is why water and oil do not mix  Polarity and molecular shape In order to predict whether a molecule will be polar or nonpolar, we must know its three-dimensional geometry CO2 and H2O WATCH: https://www.youtube.com/watch?v=72CQe-_PJU4 Before we can determine the molecular geometry, we need to draw the Lewis structure. This shows how the valence electrons of the central atom are distributed as bonds and lone pairs.  Summary for Revision Electronegativity is related to the power with which an atom in a covalent bond draw the shared bonding electrons towards it. Electronegativity increases from left to right and decreases from top to bottom in the Periodic Table. It follows a similar trend to first ionization energy, which we have seen before. You need to memorize the electronegativities of H, C, N, O, F, Cl, Br, and I. They are useful in helping you decide on the polarity of molecules. Polar molecules have partial positive and negative charges on certain atoms. The difference in electronegativity between two bonded atoms can give the percent ionic character in that bond. All polar bonds have a degree of ionic character. All ionic bonds have a degree of covalent character. Molecules can contain polar bonds, but be nonpolar overall. The crucial factor that determines whether a molecule is polar or not is the molecular geometry (shape) of the molecule. Molecular polarity affects to a large degree chemical reactivity. When a molecule is polar, it has a finite dipole moment. A nonpolar molecule has a dipole moment of zero. Polar molecules dissolve in polar solvents. Non polar molecules dissolve in nonpolar solvents. Polar liquids do not mix with nonpolar liquids (e.g. water and oil). The first step in determining the molecular geometry is to draw correctly the Lewis structure. This structure shows how the valence electrons are distributed as covalent bonds or lone pairs. We concentrate particularly on the central atom in a molecule.