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

Weird, huh

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.

Why!? Why!? Why!?

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

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

Sophia in 30 years? Must have been a rough life

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

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

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.

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.

What does that mean?!?

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.

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.

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.

Why you wear sunscreen

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 2s 2 3p 2 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 2s 1 3p 3 and now we have four single electrons to bond with hydrogen and an sp hybrid orbital

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

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.

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

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

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

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

Covalent Bonding is kinda hairy

Seriously