Chm 2045 Exam 2 Review Ch 9 And 10

Front 1

Ionic Bonds

Back 1

metal + nonmetal
- relatively strong
- electrons transferred from one to another
- large array
example: table salt
-cations and anions form and are held together electrostatically

Front 2

Properties of Ionic Compounds

Back 2

1) Crystalline solids (made of ions)
2) High melting and boiling points
3) Conduct electricity when melted or dissolved in polar liquid
4) Many soluble in water but not in nonpolar liquid
- water molecules attract ions to dissolve compound
5) ions held in place by lattice

Front 3

Covalent Bonds

Back 3

nonmetal + nonmetal
- electrons shared and localized
- travel back and forth
example: water

Front 4

Properties of Covalent Compounds

Back 4

1) Gases, liquids, or solids (made of molecules)
2) Low melting and boiling points
3) Poor electrical conductors in all phases
4) Many soluble in nonpolar liquids but not in water

Front 5

Metallic Bonds

Back 5

metal + metal
- electrons through entire crystal lattice
- electron pool
example: sodium

Front 6

Properties of Metallic Compounds

Back 6

1) High melting and boiling points
2) conduct electricity and heat
3) hard and dense
4) malleable (can be hammered into shapes)
5) ductile (can be made into wires)

Front 7

Why do chemical bonds form?

Back 7

- they lower potential energy between the charged particles in the constituent atoms
- potential energy of bonded atoms is less than that of separate atoms

Front 8

What are the guidelines for drawing a Lewis Structure for an atom or molecule?

Back 8

1) find total # of valence electrons
2) draw backbone structure
3) ensure each atom has its duet or octet, keep in mind exceptions (row 3, 2, or 1 with empty d orbitals).

Front 9

Lewis Bonding Theory

Back 9

- atoms bond to obtain noble gas configuration if possible
- cations have Lewis symbols without valence electrons because they gave them away
- anions have Lewis symbols with 8 valence electrons
*brackets and charge around symbol
- structures used as predictors of how atoms combine to form compounds and molecules

Front 10

Lattice Energy

Back 10

- the energy required to separate a mole of an ionic solid into gaseous ions (or bond strength)
- released as heat
- gaseous cations reacting with gaseous anions to form solid compound
- electron affinity and ionization energy are also involved

Front 11

Trends in Lattice Energy

Back 11

- size: shorter distance between ions means higher stability
- charge: larger charges, more stable
- Coulomb's Law: multiply charges and divide by distance

Front 12

Born-Haber Cycle Definition and Steps

Back 12

to calculate lattice energy

Front 13

Lattice Energy and Physical Properties

Back 13

- the higher the lattice energy, the less soluble a compound is in water (harder to pull apart)
- higher lattice energies associated with higher charges on the ions
- the higher the lattice energy, higher the melting point
- halides: melting points get lower down the column

Front 14

Electronegativity and Trends

Back 14

- ability of an atom to attract electrons to itself resulting in ionic and polar covalent bonds
- F is most electronegative (x=4.0)
- O is second most electronegative (x=3.5)
- increases across table
- decreases down table
- Cs is most electropositive (x=0.7) (least electronegative)

Front 15

Bond Polarity Classification

Back 15

- dipole moment will point towards more electronegative end with a plus on more electropositive end

Front 16

Percent Ionic Character

Back 16

- 6.2D is the ideal total charge-transferred dipole moment (M ionic)

Front 17

Resonance Structures

Back 17

- resonance structures differ in the position of multiple bonds and non bonding electron. The placement of atoms and single bonds always stays the same.
- only electrons move, not nuclei
- only pi electrons, single unpaired electrons, and lone pair electrons can move
- total number of electrons in the molecule do not change and neither do the number of paired and unpaired electrons
*favored:
1) negative formal charge on most electronegative atom
2) fewer and smaller formal charges

Front 18

Formal Charge

Back 18

- when atoms end up with more or fewer electrons than they brought
- sum of formal charges in a molecule = 0
- in an ion, total = the charge
Steps
1) for each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
2) subtract that from the number of valence electrons for that atom: the difference is its formal charge

Front 19

Bond Energies

Back 19

- breaking is endothermic, takes energy (+)
- forming is exothermic, releases energy (-)
- ^Hrxn reaction energy

Front 20

Bond Lengths

Back 20

- distance between nuclei of bonded atoms
- avg bond length: b/c actual bond length depends on other atoms around the bond
- single bods are longest and weakest
- triple bonds are shortest and strongest

Front 21

VSEPR

Back 21

- valence shell electron pair repulsion
* depends on how many electron groups are bonding and how many are lone pairs
- electron groups repel one another through Coulombic forces
- repulsions between electron groups on interior toms determine geometry of a molecule
- preffered geometry: in which electron groups have max separation

Front 22

Molecular and Electron Geometry

Back 22

Front 23

Bond Polarity

Back 23

- dipole moments are vectors and add according to vector algebra (can cancel or reinforce each other)
- vectors have both magnitude and directionn

Front 24

Predicting Polarity

Back 24

1) draw lewis structure
2) determine if bonds are polar (atoms have different electronegativities)
3) determine whether polar bonds add together to give a net dipole moment

Front 25

Polarity and Solubility

Back 25

- polar attracts polar
- polar dissolves well in water (polar) and in ionic compounds
- some molecule are partially polar and nonpolar