Unit 3 – Bonding Review

Octet Rule!

8 is the magic number! Most atoms want a complete valence shell with 8 electrons

Exceptions: For elements like hydrogen, helium, lithium etc, they can be happy with 2 electrons which completely fill the first principle energy level

Atoms can gain or lose electrons OR share electrons in order to get this complete Octet

Electronegativity is a measure of the tendency of an atom to attract electrons towards itself. The higher the electronegativity, the more that atom attracts electrons! Electronegativity increases from left to right on the Periodic Table

Electronegativity decreases from top to bottom on the Periodic Table

**Fluorine is the most electronegative element**

Why is this important to know?!1. Electrons with higher electronegativity tend to gain electrons to obtain

stable octet/complete valence shell2. Electrons with lower electronegativity tend to lose electrons to obtain

stable octet/complete valence shell3. Non-metals have higher electronegativities than metals; explains ionic

bonding

Ionization energy is the energy to remove one valence electron! Metals have low ionization energies Non-metal have high ionization energies

Ion Formation and PropertiesMetals Non-MetalsCation Anion

Low electronegativity Lose electrons Form positive (+) ions Ion size is smaller than atom size

High electronegativity Gain electrons Form negative (-) ions Ion size is larger than atom size

Ion Formation:1. Group 1 – Always forms +1 ions2. Group 2 – Always forms +2 ions3. Group 17 – Always forms -1 ions4. Oxygen/Sulfur – Forms -2 ion; Nitrogen forms -3 ion5. Non-Metals – Use the top charge/oxidation state on your periodic table6. Transition Metals – Can form ions with multiple charges. Can only determine

charge in a compound

**When an atom gains or loses electrons, the resulting electron configuration is typically the same configuration as a neutral noble gas, the stable Group 18

elements**

Dot Structures – Atoms and Ions1. Draw the symbol for the atom2. Count how many valence electrons the atom has3. Begin adding electrons (represented by dots) to the top, right, bottom and

left side of the symbol4. After you have added four dots or electrons, begin to pair them

Examples:

Energetics of Bonding!1. Breaking a bond always takes energy – endothermic2. Making a bond always released energy – exothermic3. Systems tends to move towards more stable, lower energy states4. More on this when we get to kinetics and thermodynamics!

Ionic Bonding Ionic bonding involves a transfer of electrons Electronegativity difference > 1.7 Properties:

o High melting point/boiling pointo Hardo Brittleo Electrolytes (ions dissociate and can conduct electricity when

dissolved in water or another polar solvent)

Writing Ionic FormulaeIonic compounds are held together by electrostatic forces, the force of attraction between the + and – charged ions (opposite charges attract!). The overall resulting compound is neutral. In order to write the formula for a compound, you must balance the charges!

**Use the Kriss Kross method!!**

Remember: Ionic formulas are ratios of the anions and cations in the salt crystal. If you can reduce, you must!

Polyatomic Ions

Polyatomic ions are compounds where the individual atoms are covalently bonding but with an overall charge. They can therefore enter into ionic compounds with ions of the opposite charge. Polyatomics ions are found in your reference table!

Dot Structures – Ionic Bonding

Covalent BondingA covalent bond is formed when two non-metals share electrons in order to form a stable octet

Lewis Dot Structures Used to represent bonding among covalent compounds Show which atoms are bound together and how many bonds are present

Steps To Draw:1. Determine # and types of atoms in compound2. Count total number of valence electrons. This is the number you must

show in your picture!3. Arrange skeleton and use single bonds (a line) to connect all atoms

a. Carbon, if present, is usually in the centerb. Hydrogen and the halides (Group 17 can never be in the center)c. Otherwise, least electronegative atom in center

4. Start to add lone pair electrons.5. If you cannot make each atom happy, i.e. give it a stable octet, you need to

start adding double and triple bonds6. Count you electrons! They should add to the original # of electrons you

started with!!

Polar vs. Non-Polar BondsIf the electronegativity difference between two atoms bound together is >0.4 and <1.7, the electrons will be shared unevenly between them and the bond will be polar.

The more electronegative atom gets a partially negative sign as the electrons spend more time around its nucleus

The less electronegative atom gets a partially positive sign as the electrons spend less time around its nucleus

Polar vs. Non-Polar CompoundsThe polarity of a compound is determined by its structure. If the compound is symmetrical - i.e. it has either a linear or tetrahedral shape AND the same atoms all bound to the central atom - the overall compound EVEN though it might have polar bonds can be non-polar. This is because the forces that are pulling the electrons cancel out. An example of this would be CO2 or CCl4. In the later, the C-Cl BOND is polar but the overall compound CCl4 is non-polar because it is symmetrical and the forces pulling on the electrons cancel out.

Molecular Geometry

VSEPR theory – electrons clouds want to be as far apart from each other as possible. Effects the overall shape of molecules.

# Atoms Bonded To Central Atom

# Of Lone Pairs Bonded To Central Atom

Shape Bond Angles

2 0 Linear 1802 1 Bent 1202 2 Bent 109.53 0 Trigonal Planar 1203 1 Trigonal Pyramidal 109.54 0 Tetrahedral 109.5

Nomenclature – Ionic and Covalent Compounds

Ionic Covalent1. First, state the name of the cation

(including polyatomic cations such as ammonium)

2. ONLY IF the cation is a transition metal, use a Roman Numeral to write the positive charge on the anion (e.g. FeO = Iron (II) Oxide)

3. Finally, state the name of the anion with its ending changed to “-ide” if it is a non-metal element. If it is a polyatomic ion, keep its full name and original ending (-ide, -ate or –ite)

1. Write the prefix and name of the least electronegative non-metal. Exception: if the least electronegative metal has a subscript 1, do not use the prefix mono- for the first element named.

2. Write the prefix (including mono- if the subscript is only 1) of the second, more electronegative non-metal and change the ending to –ide

H20 = dihydrogen monoxide

Metallic Bonds Between one or more metals Metals have few valence electrons and low electronegativities and ionization

energies Metal atom is considered to have a central portion on “kernel” made of

nucleus and on valence electrons Positively (+) charged kernals are held together in a “sea of electrons” –

valence electrons freely float among these kernels in the metallic crystal Properties:

o Ductileo Malleableo Hardo High melting pointso High boiling points

o Conduct heat and electricity well (even in solid state)

Coordinate Covalent CompoundsA coordinate covalent compound forms when one atom contributes both electrons in the bond to the other atom.

Example: NH3 + H+ ------------ NH4+1

Network SolidsNetwork solids are covalent compounds where a large number of atoms are all bonding together. These extremely large covalent compounds are extremely hard. The classic two examples are diamond (a large number of carbon atoms all bound together) and –(SiO2)- (quartz).

AllotropesWhen an element can exist in more than one form, the multiple forms are called allotropes.

Example: Carbon can be amorphous carbon, graphite or diamond; oxygen can be the kind we breathe (O2) or ozone (O3)

Properties of Ionic versus Covalent Bonds!

Ionic solids hard brittle high melting point form electrolyte solutions in water (solutions that conduct electricity) form crystals with regular geometric shapes that cleave along cleavage

planes - saltCovalent substances

SolidsNetwork solids - a large macromolecule held together by covalent bonds

Hard Brittle High melting point Don't conduct electricity solid or melted Form crystals with regular geometric shapes that cleave along cleavage

planes - diamonds (C), sand (SiO2)

Molecular solids - discrete molecules held together by intermolecular forces May be soft May have a low melting point Don't conduct electricity solid or melted

Liquids Polar compounds dissolve best in other polar compounds (acids and water) Nonpolar compounds are often immiscible in water and dissolve better in

nonpolar compounds (tar in gasoline)

Intermolecular ForcesIntermolecular forces are the forces BETWEEN molecules that hold them together. These forces are much weaker than the strength of covalent and ionic bonds. They are all due to the attraction of partial positive charges and partial negative charges in a molecule because attracted to the opposite partial sign in another molecule.

Three types of intermolecular forces can operate between covalent molecules:

1. Dispersion Forces - also known as London Forces (named after Fritz London who first described these forces theoretically 1930) or as Weak Intermolecular Forces or as van der Waal's Forces (named after the person who contributed to our understanding of non-ideal gas behavior).

2. Dipole-dipole interactions – between two polar molecules3. Hydrogen bonds – compounds that have F, O or N bonded to H

A fourth intermolecular force you should know is ion-dipole (also known as molecule-ion)! This is when an ionic compound is dissolved into a polar substance, such as water, and the ionic crystal dissolved.

4. Ion-Dipolar (Molecule-Dipole) – when an ionic compound dissolves in a polar covalent substance. + ions attracted to partially negative side of molecule and – ions attracted to partially positive side of molecule. Example: NaCl in water.

Relative strength of Intermolecular Forces:

dispersion forces < dipole-dipole interactions < hydrogen bonds < ion-dipole

Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's Forces)

VERY WEAK forces of attraction between molecules resulting from a momentary dipoles occurring due to uneven electron distributions in neighboring molecules as they approach one another (imagine the electron cloud temporarily behind concentrated mostly on one side of the molecule and therefore temporarily induced a partial negative and partial positive charge)

Dispersion forces are the only type of intermolecular force operating between non-polar molecules!!!!!!!!!

As the mass of the molecules increases, so does the strength of the dispersion force acting between the molecules, so more energy is required to weaken the attraction between the molecules resulting in higher boiling points

Dipole-dipole Interactions are stronger intermolecular forces than Dispersion forces occur between molecules that have permanent net dipoles (polar molecules),

for example, dipole-dipole interactions occur between SCl2 molecules, PCl3 molecules and CH3Cl molecules.

The partial positive charge on one molecule is electrostatically attracted to the partial negative charge on a neighboring molecule.

Hydrogen bonds occur between molecules that have a permanent net dipole resulting from

hydrogen being covalently bonded to either fluorine, oxygen or nitrogen. Strongest type of intermolecular force between molecules (ion-dipole is

stronger but contains ions) Important in biology!!! Holds your DNA together! Allows you to

breathe! Important in biology!!! Explains high boiling point of water!

H-bonds are stronger intermolecular force than either Dispersion forces or dipole-dipole interactions since the hydrogen nucleus is extremely small and positively charged and fluorine, oxygen and nitrogen being very electronegative so that the electron on the hydrogen atom is strongly attracted to the fluorine, oxygen or nitrogen atom, leaving a highly localized positive charge on the hydrogen atom and highly negative localized charge on the fluorine, oxygen or nitrogen atom. This means the electrostatic attraction between these molecules will be greater than for the polar molecules that do not have hydrogen covalently bonded to fluorine, oxygen or nitrogen.

Effect of Intermolecular forces on melting and boiling points of molecular covalent substances:

Explains extremely high boiling and melting points of H2O, NH3, etc.