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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
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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
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.
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
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.
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.
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
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.