Covalent Bonds! Yeah!

1. Covalent Bonds! Yeah! • Elements with high electronegativities (non-metals) will not give up electrons. Bonds are not formed by a transfer of electrons, they are formed by sharing electrons. • Molecules are neutral groups of covalently bonded atoms • A diatomic molecule is two atoms of the same element covalently bonded together

2. Weird, huh

3. Molecular Compounds • Molecular compounds tend to have lower melting points than ionic compounds • Many of them are either gases or liquids at room temp. • Some molecules can conduct electricity but most don’t. • Polyatomic ions are covalently bonded atoms with a charge.

4. Why!? Why!? Why!?

5. Octet Rule…. again • Atoms what to attain the electron configuration of a noble gas, (8 electrons in the outer shell) • Nonmetal will share from 1-3 electrons in order to achieve eight. • Single covalent bonds, two shared electrons • Double, four shared. Triple, 6 shared • Each shared pair makes a bond

6. More sharing • Some electrons will not be involved in the bonding process and are called an unshaired pair. Single electrons are always bonded. • A dashed line represents a bond, multiple dashes, multiple bonds. • Some molecules are exceptions to the octet rule, multiple bonds make up for this, NO2 • Chemical symbols with dashes represent a structural formula, compared to a chemical formula which is just symbols and subscripts

7. Sophia in 30 years?Must have been a rough life

8. Polyatomic ions • Covalently bonded atoms with a charge, several of them • Many ionic compounds end either “ate” or “ite” • Many of them are coordinate covalent compounds. • Coordinate covalent compounds are compounds where one atom donates both bonding electrons. NH3 and NH4 for example

9. Resonance • Resonance structures are different electron dot configurations for the same molecule • Ozone, for example can be drawn 2 different ways.

10. Bond Dissociation Energy • The energy required to break a covalent bond. • A large bond dissociation energy corresponds to a strong covalent bond. • Single bond is weaker than a double weaker than a triple. • Some single bonds can be stronger than other single bonds.

11. Molecular Orbitals • At0ms have atomic orbitals. When atoms bond together, it is theorized that these orbitals overlap to form molecular orbitals, or a combination of the two atomic orbitals. • A sigma bond forms when two orbitals are symmetrical around the two nuclei or the axis between them , s or p orbitals for example • A pi bond forms when p orbitals overlap side by side, electrons are found above and below the bonding axis.

12. What does that mean?!?

13. Molecular Orbitals • Pi bonds overlap less than sigma bonds and are weaker than sigma bonds • This is one of several theories to explain the principles behind atomic bonding, how it occurs, and the shapes that result.

14. VSEPR • Valence Shell Electron Pair Repulsion Theory, notice it says theory. Another way to try to explain molecular bonding. • According to this theory, valence shell electron pairs repel each other in order to stay as far apart as possible. • This accounts for bonding electrons and unbonded pairs.

15. VSEPR • Shapes include Linear triatomic, trigonal planar, bent triatomic, pyramidal, and many others. The shape depends on the number of atoms, bonds, and unbonded electrons.

16. Why you wear sunscreen

17. Hybridization • Long story short, different orbitals in the same atom form one hybrid orbital in that atom • Methane, CH4 for example, Carbon has an outer configuration of 2s2 3p2 It has to bond with four hydrogens, but there are only 2 unpaired electrons. One electron comes up from the s orbital to the p orbital to make it 2s1 3p3 and now we have four single electrons to bond with hydrogen and an sp hybrid orbital

18. Polar Bonds • Covalently bonded atoms become polar when one atom has a higher electronegativity than the other. (usually, just more electrons) • A polar covalent bond is one where atoms are shared unequally. One side of the molecule develops a positive charge and the other side develops a negative charge due to the imbalance of electrons

19. Polar • Polar covalent bonds form polar molecules • Polar bonds can cancel each other out if they are in the same plane and linear, CO2 for example • Polar molecules are attracted to each other by opposite charges.

20. If you are watching from the ski lodge below, you might think about moving

21. Intermolecular forces • Molecules are attracted to each other by a variety of ways called intermolecular forces. • Intermolecular forces are weaker than atomic forces such as covalent or ionic bonds. • The two weakest forces are collectively called Van der Waals Forces. They are dipole and dispersion. • Dipole is the same as polar, the negative end of one molecule is attracted to the positive end of another

22. More intermolecular • After dipole are dispersion forces , the weakest of all intermolecular forces. • Dispersion is due to the movement of electrons and is slightly stronger with more electrons present. • Hydrogen bonds, the strongest, occur between molecules that due to their polarity, share a hydrogen, same as polar or dipole but with a hydrogen in the middle

23. Hydrogen Bonds • They are the strongest and account for a lot of important properties in water and biological processes.

24. Covalent Bonding is kinda hairy

25. Seriously