Covalent Compounds

1. Covalent Compounds Covalent Bonds, Drawing and Naming Molecules, and Molecular Shapes

2. Covalent Bonds • Sharing Electrons! • Electrons rearranged when an ionic bond forms, electrons transfer from one atom to another to form charged ions. • In covalent bonds, neutral atoms SHARE electrons. • Do not make ions (cations or anions), they share so everyone is happy.

3. Forming Molecular Orbitals • Both nuclei repel each other, as do both electron clouds. • The nucleus of each atoms attracts both electron clouds. • These two smash the atoms together and form a single molecular orbital bound via a covalent bond. • Molecular Orbital – The space that these shared electrons move within (high probability)

4. Energy and Stability • Most individual atoms are highly unstable, with the exception of _____________________. • Unbonded atoms have a high potential energy. This stored energy is released when they form compounds. • When diatomic hydrogen forms a compound, energy is released to the surroundings. Now a high stability and a low potential energy.

5. Energy is Released • As two hydrogens come near one another, energy is released as the attractive forces pulls the atoms closer. • Eventually they get close enough where the attracting forcing of the electrons and nucleus are balanced out by the repelling forces of the nucleus v. nucleus and elections v. electrons. • Once bonded, the atoms are at their lowest potential energy. Any closer together and there will be repelling forces acting causing them to want to pull apart.

6. Potential Energy and Bond Length • The distance between two bonded atoms at their minimum potential energy is known as the bond length. • However, the two nuclei are in constant motion. They vibrate back and forth, pushing and pulling. • The bond length is an average distance to account for the vibrations.

7. Energy and Bond Length • At the bond length of 75 pm, the potential energy of H2 is -436 kJ/mol. • This means that 436 kJ of energy is released when 1 mol of H2 is formed. • It also means that 436 kJ if energy must be supplied to break the bonds and separate the hydrogen atoms in 1 mol of H2 molecules. • The energy required to break a bond between two atoms is the bond energy. • The stronger the bonds, the shorter the bond length.

8. Electronegativity • Two hydrogens attract and repel one another equally because they have the same electronegativity. • Many other atoms with different electronegativities form covalent bonds. • A nonpolar covalent bond is a bond in which the bonding electrons in the molecular orbital are shared equally. • A polar covalent bond is a covalent bond in which a shared pair of electrons is held more closely by one of the atoms.

9. Polar Covalent Bonds • In a polar covalent bond, the shared electrons, which are in a molecular orbital, are more likely to be found nearer to the atoms whose electronegativity is higher. • If the electronegativity of one of them is high enough, then the atom may PULL off an electron forming a ______________________ bond.

10. Electronegativities • The electronegativity difference between oxygen and magnesium is high enough to allow oxygen to PULL off two electrons from magnesium, forming an ionic compound. • You can classify bonds based on the difference between electronegativities.

11. Electronegativities • Even electron distribution is a _____________________. • This bond occurs at an electronegativity difference of 0.5 or lower. • Uneven electron distribution is a __________________. • This occurs between 0.5 and 2.1 • Separate electron clouds is a(n) ____________ bond. • This occurs between 2.1 and 3.3 **These are arbitrary values and not exact** • Also, remember that covalent bonds usually occur between two ____________________.

12. Positive and negative Ends • When there is an uneven electron distribution as in _________________ bonds, there is a more positive side and a more negative side. • Since the electrons are “more attracted” to one side, that side is more negative. We call this “partially negative” denoted δ- • A molecule or part of a molecule that contains both positively and negatively charged regions is known as a dipole.

13. Careful, Not Ionic! • This δ+ and δ- region of a polar covalent molecule does not mean positive and negative as in an ionic bond. Electrons are not stripped away, but simply shared. • One atom just WANTS the electron(s) more and thus is a little more negative than the other side (δ-)

14. Polarity and Bond Strengths • The greater the difference between the electronegativity values of two elements joined by a bond, the greater the polarity of the bond. • The greater the electronegativity differences tend to be associated with stronger bonds.

15. Bond Character • We know that Na and F form a _____________ bond. • The electronegativity difference is 3.1 • Ca and O form a ______________ bond. • The electronegativity difference is 2.4 • Based on these values, we can say that Na and F have a bond between them with a higher percentage of ionic character.

16. Bond Character • The same is true with polarity • A bond between C and Cl has a difference of 0.6 • A bond between Al and Cl has a difference of 1.6 • This information tells us that Al and Cl will form a MORE polar bond than C and Cl

17. Chemical and Physical Properties • The type of bond that forms determines the physical and chemical properties of the substance. • Consider solid potassium (bunches of K together) • Metal bonds are the result of the valence electrons being attracted to all of the atoms in the solid. • These valence electrons can move easily from one atom to another. • They are free to roam around in the solid and can conduct electric current (a physical property).

18. Test Coming up: Ionic bonding and covalent bonding (electronegativities)