Thursday 10.27.16 Agenda Review POGIL exercise Do Now Covalent Bonding notes LDD for compounds Do Now Pull out your POGIL packets and a scrap sheet of paper Homework

Covalent Bonding

Bonding Two methods to gain or lose valence electrons: Transfer of Electrons = Ionic Bonding Sharing of Electrons = Covalent Bonding Ionic Bonding - attracted to each other, but not fully committed Covalent Bonding - fully committed, and shares everything

Why do metals have multiple charges? Transition metals have 2 s electrons in their valence shells They also have ‘high energy’ d electrons These atoms typically lose their s electrons to form a +2 ion, but also sometimes lose some of their d electrons This is why there are two possible oxidation states for many transition element

Naming Covalent Compounds Prefixes used to denote number of atoms of element The suffix -ide is added to the more electronegative element For the least electronegative (on the left), mono- is not used to denote one atom. The first vowel is often dropped to avoid the combination of “ao” or “oo”. CO = carbon monoxide (monooxide)

Molecules Molecule = covalently bonded atoms Diatomic molecule = two atoms Molecular compounds – lower melting and boiling point than ionic compounds Molecular formula Doesn’t have to be lowest whole-number ratio Does represent structure

Molecular Representations

Lewis dot structures of covalent compounds In covalent compounds atoms share electrons. We can use Lewis structures to help visualize the molecules. Lewis structures Multiple bonds must be considered. Will help determine molecular geometry. Will help explain polyatomic ions.

Types of electrons H Cl Bonding pairs Two electrons that are shared between two atoms. A covalent bond. Unshared pairs A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. oo Unshared pair H Cl oo oo oo Bonding pair

Covalent Bonding When two similar atoms bond, none of them wants to lose or gain electrons Share pairs of electrons to each obtain noble gas e- configuration. Each pair of shared electrons = one covalent bond Unshared Pairs = Pairs of e- not shared by all atoms Show unshared pairs as dots N H

H H H H C H F F H Single covalent bonds Do atoms (except H) have octets?

Sharing of Electrons Water - H2O - Covalent Bonding O H H = 1e-, O = 6e- Each H shares their electron with O Þ O = 8e- = [Ne] O shares 1e- with each H Þ H = 2e- = [He] Use dashes or dots to show covalent bonds H O

Multiple Covalent Bonds Elements can share more than two electrons - creates double and triple bonds Carbon Dioxide (CO2) - 4 electrons shared between the carbon and the oxygens (double bond) O C O C

Triple Covalent Bonds C H C H Triple bond = Six electrons shared between the carbon atoms Ethyne (C2H2), a.k.a acetylene Bond length decreases from single to triple bonds Bond strength increases from single to triple bonds C H C H

Coordinate Covalent Bonds Coordinate covalent bond = one atom contributes both bonding electrons Polyatomic ions contain coordinate covalent bonds O = 8e- C = 6e- C O O = 8e- C = 8e- C O O must contribute one of its pairs

Drawing Lewis structures Write the symbols for the elements in the correct structural order. Calculate the number of valence electrons for all atoms in the compound. Put a pair of electrons between each symbol, the bond between each. Beginning with the outer atoms, place pairs of electrons around atoms until each has eight (except for hydrogen). If an atom other than hydrogen has less than eight electrons, move unshared pairs to form multiple bonds.

Lewis structures C-O-O O-C-O Example CO2 Step 1 Draw any possible structures C-O-O O-C-O You may want to use lines for bonds. Each line represents 2 electrons.

Lewis structures Step 2 Determine the total number of valence electrons. CO2 1 carbon x 4 electrons = 4 2 oxygen x 6 electrons = 12 Total electrons = 16

Lewis structures C O O Step 3 Try the C-O-O structure Try to satisfy the octet rule for each atom - all electrons must be in pairs - make multiple bonds as required Try the C-O-O structure No matter what you try, there is no way satisfy the octet for all of the atoms. C O O

Lewis structures O C O O=C=O = is a double bond, This arrangement needs too many electrons. O C O How about making some double bonds? O=C=O That works! = is a double bond, the same as 4 electrons

Ammonia, NH3 H H N H H H N H Step 1 Step 2 3 e- from H 5 e- from N 8 e- total Step 3 N has octet H has 2 electrons (all it can hold) H H N H

Chapter 8

Resonance Structures Structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms = Resonance Structures Example: Ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

Resonance Structures Resonance Structures

Resonance structures O - S = O O = S - O Sometimes we can have two or more equivalent Lewis structures for a molecule. O - S = O O = S - O They both - satisfy the octet rule - have the same number of bonds - have the same types of bonds Which is right?

Resonance structures O - S = O O = S - O O S O They both are! But really…neither are. O - S = O O = S - O O S O This results in an average of 1.5 bonds between each S and O.

Exceptions to Octet Rule Three classes of exceptions to the octet rule for molecules Odd number of electrons One atom has less than an octet One atom has more than an octet Odd Number of Electrons Few examples. Molecules such as ClO2, NO, and NO2 have an odd number of electrons.

Exceptions to Octet Rule Less than an Octet Relatively rare. Typical for compounds of Groups 13. Most typical BF3. More than an Octet Atoms from the 3rd period onwards can accommodate more than an octet. The d-orbitals are low enough in energy to participate in bonding and accept extra electron density.

An example: SO42- S O 1. Write a possible 2. Total the electrons. arrangement. 2. Total the electrons. 6 from S, 4 x 6 from O add 2 for charge total = 32 3. Spread the electrons around. S O - |

Atoms with fewer than eight electrons Beryllium and boron will both form compounds where they have less than 8 electrons around them. : : :F:B:F: :F: :Cl:Be:Cl: : : : : :

Species with an odd total number of electrons Example - NO Nitrogen monoxide is an example of a compound with an odd number of electrons. It is also known as nitric oxide. It has a total of 11 valence electrons: 6 from oxygen and 5 from nitrogen. The best Lewis structure for NO is: :N::O: : .

Electronegativity Electronegativity = Ability to attract electrons in a chemical bond Decrease as you move down a group Increase as you move from left to right across a period. Decreasing Electro-negativity Increasing Electro- negativity

Polar Covalent Bonds Sharing of electrons in a covalent bond does not imply equal sharing of those electrons. In some covalent bonds - electrons located closer to one atom than the other Unequal sharing Þ polar bonds. Sharing based on electronegativity of elements in bond

Electronegativities Chapter 8

Electronegativity and Polarity Difference in electronegativity between atoms is a gauge of bond polarity Difference Type of Bond Electrons 0 – 0.4 Nonpolar Covalent Equal sharing 0.4 – 1.0 Moderately Polar Covalent Unequal sharing 1.0 – 2.0 Very Polar Covalent » 3 Ionic Transfer

Nonpolar and polar covalent bonds When two atoms share a pair of electrons equally. H H Cl Cl Polar A covalent bond in which the electron pair in not shared equally. H Cl Note: A line can be used to represent a shared pair of electrons. oo oo oo oo oo oo oo oo oo d+ d- oo oo oo

Polar Dipoles O H δ- δ+ δ+ The positive end (or pole) in a polar bond is represented + and the negative pole - = dipoles (partial charges) Polar molecules placed between electric (+/-) plates – become aligned with plates δ- H O δ+ δ+

Attractions between Molecules Intermolecular attractions weaker than ionic or covalent bonds these are attractions BETWEEN molecules not inside molecules like ionic or covalent bonds (intramolecular)

Types of Intermolecular Attractions van der Waals forces Dipole interaction δ+ end of one molecule attracted to δ- end of another Dispersion forces Caused by movements of e- Hydrogen Bonds Hydrogen bonded with very electronegative element is also weakly bonded to an unshared pair on another electronegative atom

Characteristics of Ionic and Covalent Compounds Ionic Compound Covalent Compound Representative Unit Formula Unit Molecule Bond Formation Transfer of e- Sharing of e- Type of elements Metal + Nonmetal Nonmetals Physical State Solid Solid, liquid, or gas Melting Point High (usually above 300°C) Low (usually below 300°C) Solubility in water High High to low Electrical conductivity of aqueous solution Good conductor Poor to nonconducting