U2: Molecular and Ionic Compound Structure and Properties

Types of Bonds

● Bonds form in order to decrease potential energy and increase stability

● Chemical Bond:

○ When an attractive force between atoms or ions binds them together as a unit.

● Covalent Bonding (AKA Molecular Bonding)

○ When two nonmetals share electrons

○ Nonpolar Covalent (Equal Sharing) and Polar Covalent (Unequal Sharing)

● Metallic Bonding

○ Metals are held together by sharing their outer electrons.

○ Those ecan flow around atom to atom, which is represented by the term: sea of electrons.

● Ionic Bonding

○ Metal and nonmetal transfer electrons

○ Very Polar

○ Lattice Energy = the energy that is released when a crystal forms from ions

■ Delta HLE = (Q₁Q₂ / r)

● Q = charge of ions

● r = distance between nuclei

Ion Formation - Ionic Compounds

● Ions are atoms which gain or lose electrons

● Ions have overall charge

○ Noble Gas Rule: Main group elements want to have the same number of electrons as the closest noble gas. They gain or lose electrons in order to do this

■ Losing Electrons = Positive Ions = cations

■ Gaining Electrons = Negative Ions = anions

■ Ex. Sodium: 11 protons and 11 electrons

● Closest noble gas: Neon (10e-)

● Sodium loses 1 eto have the same number as neon: Na+1, Sodium Ion

Naming Ions

● Positive Ion = Name doesn’t change

● Negative Ion = Name’s ending changes to -ide

Writing Ionic Formulas

● Positive Ions only bond to negative ions

● Positive Ion is always first

● Positive Charge must balance the negative

● Subscripts are used after each symbol to indicate how many are needed (unless it’s a 1)

Naming Ionic Formulas

● Name the Cation (+), Anion (-)

● The name shouldn’t include how many atoms of each ion are present

● Examples

○ Na+1 Br-1 = NaBr, Sodium Bromide

Transition Metals

● Most transition metals can have more than one charge

● Specify the charge in the name with Roman Numerals

○ Fe+2 is Iron (ll)

● Exceptions: There are three transition metals that only have one possible charge

○ Zn = +2, Cd = +2, Ag = +1

○ For these ions, we don’t use roman numerals in the name

Polyatomic Ions

● Clusters of atoms that have an overall charge

● They are attracted to opposite ions and form compounds

● When using polyatomics in compounds, NEVER alter their formulas

● If more than one needed, use parentheses around the formula followed by the subscript

● List of common polyatomic ions that you should memorized or be familiarized with:

Molecular Compounds

● Made of all nonmetals that share electrons

● Use prefixes to indicate how many atoms of each element are in the middle

○ Mono - (1)

○ Di - (2)

○ Tri - (3)

○ Tetra - (4)

○ Penta - (5)

○ Hexa - (6)

○ Hepta - (7)

○ Octa - (8)

○ Nona - (9)

○ Deca - (10)

● When there is only one atom of the first element, leave the “mono” off

● Second element ends with -ide

Polarity

● Polar bonds occur when e- are shared unequally

● A molecule is polar if it has polar bonds that are not symmetrical

● A polar molecule has one side that is positive and the other side is negative

Diatomic Molecules

● Composed of 2 nonmetal atoms

● The H-7 Club

● There are 7 and they make a 7 on the periodic table, which starts at #7

○ H2, N2, O2, F2, Cl2, Br2, L2

● Examples

○ CCl4 = Tetrachloride

○ N2O3 = Dinitrogen Trioxide

○ P2F5 = Diphosphorus Pentafluoride

Hydrates

● A solid compound that containing or linked to water molecules

● Name the ionic compound followed by the numerical prefix and the suffix -hydrate

○ CuSo4 * 5 H2O = Copper (ll) Sulfate Penta Hydrate

Lewis Diagrams

● Representation of the valence electrons in a molecule

● Show how electrons are arranged around individual atoms in a molecule

● The dots represent the valence electrons and the line is shown for a bonding of electrons

between two atoms

○ 1 line = Single Bond = 2 electrons

○ 2 lines = Double Bond = 4 electrons

○ 3 lines = Triple Bond = 6 electrons

○ 2 dots = lone pair of electrons

● Steps for drawing a Lewis Diagram

○ Count all the valence electrons

○ Determine the central atom (the element there is only one of)

○ Put all the remaining valence electrons on atoms as lone pairs

○ Turn lone pairs into double/triple bonds to give every atom an octet (or duet)

● Alternate method to drawing a Lewis Structure

○ Figure out how to many bonds each atom wants

○ Put the atom that wants the most bonds in the middle

○ Give each atom the number of bonds it wants

○ Add lone pairs to give every atom an octet

Valence Electron Shell Pair Repulsion (VSEPR)

● VSEPR predicts shapes of molecules

● Can be found by using the number of electron pairs

○ To predict the shape, first draw out the Lewis Structure of the compound and look at its central atom

○ Count its valence electrons

○ Add one electron for each bonding atom

○ Add or subtract electrons for charge

○ Divide the total of these by 2 to find the total number of electron pairs

○ Use this final number to predict the shape

● Example

○ (PF6)-

■ Central Atom = Phosphorus

■ Valence Electrons on Central Atom = 5

■ 6 (F) Atoms = 6

■ +1 Electron for negative charge on P = 1

■ Total = 12, divide by 2 to become 6

■ Phosphorus is an octahedral

Hybrid Orbitals

● The central atom’s coordination number gives the number of hybrid orbitals

● If any element has a coordination number of 4, it will hybridize all four valence orbitals (sp3)

● If an element has a coordination number of 3, it only needs to hybridize three of its orbitals (1 remains as a normal “p” orbital), it is (sp2)

● Hybrid orbitals make sigma bonds, which is the direct overlap of orbitals between two atoms

● Non-hybridized orbitals can still bond but they are different than the hybrid orbital ones 

● Non-hybrid orbitals make pi bonds

○ Sharing of electrons between parallel p orbitals