Chemistry Test #5

Front 1

Coulomb's Law

Back 1

1. For like charges, the potential energy is positive and decreases as particles get farther apart!

2. For opposite charges, the potential energy is negative and becomes more negative (increases) as get closer together.

Front 2

What type of bond - metal and nonmetal?

Back 2

1. Ionic (electrons transferred)

Front 3

What type of bond - nonmetal/nonmetal?

Back 3

1. Covalent (electrons shared)

Front 4

What type of bond - Metal/Metal?

Back 4

1. Metallic (electrons pooled)

Front 5

Why is Helium an exception to the Lewis Dot Structure?

Back 5

1. It fills and is stable at 1s2, which would be only two dots (duet).

Front 6

In Lewis' theory, a chemical bond is__________.

Back 6

1. sharing or transferring of electrons to attain stable electron configurations for bonding atoms.

Front 7

When you combine ionic compounds, such as K and Cl in Lewis dot structures, you get two ________.

Back 7

1. Ions, K+ and [Cl]-
2. The Lewis structure shows KCl connected though.

Front 8

What is Lattice Energy?

Back 8

1. energy associated with forming a crystalline lattice from gaseous ions.

Front 9

Born-Haber Cycle...

Back 9

1. hypothetical series of steps that represent the formation of an ionic compound from its constituent elements.

Front 10

Trends in Lattice Energy

Back 10

1. Lattice Energy becomes less negative with increasing ionic radius.
2. Lattice Energy becomes more negative with increasing magnitude of ionic charge.
3. Lattice energy decreases as you go down column.

Front 11

Do ionic compounds tend to conduct electricity as solids? How about as liquids?

Back 11

1. DO NOT conduct as solids
2. DO conduct as liquids.

Front 12

Bonding pair...

Back 12

1. Pair of electrons shared between two atoms.

Front 13

Lone pair...

Back 13

1. Pair of electrons that is associated with only one atom.

Front 14

Which bond has the shortest length and most strength?

Back 14

1. Triple
2. Order goes Triple, Double, Single in terms of shortest to longest length of bond.

Front 15

Covalent bonds are highly__________ compared to ionic bonds which are ______________.

Back 15

1. Directional
2. Nondirectional

Front 16

Polar Covalent Bond

Back 16

1. Covalent bond with significantly different electronegativities, resulting in an unequal distribution of electron density.

Front 17

Electronegativity...

Back 17

1. Ability of an atom to attract electrons to itself in a chemical bond.

Front 18

Trends in electronegativity...

Back 18

1. INCREASES across a row/period
2. DECREASES down a column
3. Flourine is most electronegative.
4. Francium is least electronegative.

Front 19

The degree of polarity in a chemical bond depends on _____________________.

Back 19

1. Electronegativity difference between the two bonding elements.
2. Greater the electroneg. difference, the more polar the bond.

Front 20

The greater the electronegativity difference, the _______________.

Back 20

1. more polar the bond.

Front 21

Nonpolar electrons are ____________________.

Back 21

1. Electrons shared equally

Front 22

Dipole moment...

Back 22

1. A measure of separation of positive and negative charge in a molecule.
2. Measure of Bond Polarity

Front 23

In Lewis dot structures, _________________ go in the middle.

Back 23

1. Less electronegative.

Front 24

Resonance Structures

Back 24

1. Two or more valid Lewis dot structures, shown with double arrow to indicate actual structure is intermediate between them.
2. Double bond could be one either H, etc.

Front 25

Formal Charge

Back 25

1. Charge atom would have if all bonding electrons were shared equally between bonded atoms.
2. Formal Charge of H in HF = Valence electrons of H - (Number of electrons H "owns")

Front 26

Sum of all formal charges in a neutral atom must be... how about in an ion?

Back 26

1. Must be zero in neutral
2. Must be charge of ion in an ion

Front 27

Free Radicals

Back 27

1. Molecules or ions with an odd number of electrons (odd-electron species).

Front 28

Incomplete Octets...

Back 28

1. Molecules or ions with fewer than 8 electrons around an atom.
2. Boron is a good example

Front 29

Expanded Octets...

Back 29

1. molecules or ions with more than 8 electrons around an atom.
2. Elements in third row and beyond often display expanded octets.

Front 30

Bond energy

Back 30

1. Energy required to break 1 mole of a bond in a gas phase.
2. ALWAYS POSITIVE because it takes energy to break a bond.

Front 31

Compounds with stronger bonds tend to be more ___________ and less chemically reactive. Therefore, have __________ bond energy.

Back 31

1. stable
2. MORE

Front 32

A reaction is ________ when weak bonds break and strong bonds form.

Back 32

1. Exothermic

Front 33

A reaction is _______ when strong bonds break and weak bonds form.

Back 33

1. Endothermic

Front 34

Forming of bonds, __________ energy.

Back 34

1. Releases

Front 35

Bond Length

Back 35

1. Average length of a bond between two particular atoms in a large number of compounds.
2. Depend on type of atom involved and type of bond.

Front 36

Valence Shell Electron Pair Repulsion (VSEPR)

Back 36

1. electron groups repel each other through coulombic forces.

Front 37

Preferred geometry is one in which electron groups ___________________.

Back 37

1. Have maximum separation and therefore lowest energy possible.

Front 38

Molecular geometry depends on...

Back 38

1. The number of electron groups around central atom.
2. How many of these groups are non-bonding versus bonding.

Front 39

Two Electron Groups = __________.

Back 39

1. Linear Geometry
180 degree angles.

Front 40

Three Electron Groups = __________.

Back 40

1. Trigonal Planar Geometry
120 degree angles.

Front 41

Four electron groups = ___________.

Back 41

1. Tetrahedral Geometry
109.5 degree angles

Front 42

Five Electron Groups = ____________.

Back 42

1. Trigonal bipyramidal
120 degree between equatorial positions
90 degree between axial and equatorial

Front 43

Six Electron Group = ____________.

Back 43

1. Octahedral Geometry
90 degree angles

Front 44

LP-LP _____ LP-BP______BP-BP

Back 44

1. >>
2. LP-LP > LP-BP > BP-BP

Front 45

Arrow points in direction of ________ electronegative atom.

Back 45

1. MOST

Front 46

A bond is polar if the two bonding atoms______________________.

Back 46

1. Have sufficiently different electronegativities.

Front 47

If vector (arrows for electronegativity) sum to zero, the molecule is _____________.

Back 47

1. Non-polar

Front 48

Valence Bond Theory

Back 48

1. Valence electrons reside in quantum-mechanical orbitals.
2. Chemical bond results from overlap of two half-filled orbitals with spin-pairing of the two valence electrons.
3. Shape of the molecule is determined by the geometry of the overlapping orbitals.

Front 49

Hybridization

Back 49

1. Mathematical process where standard atomic orbitals are combined to form new orbitals called hybrid orbitals.

Front 50

Number of standard atomic orbitals added together equals the ___________________.

Back 50

1. Number of hybrid orbitals formed.
2. Total number of orbitals is conserved.

Front 51

Particular combinations of standard atomic orbitals added together determines the ____________________________.

Back 51

1. Shapes and energies of the hybrid orbitals formed.

Front 52

Particular type of hybridization that occurs is the one that yields the _____________________.

Back 52

1. Lowest overall energy for the molecule.

Front 53

Sigma Bond

Back 53

1. Resulting bond forms between a combination of any two s, p or hybridized orbitals that overlap end-to-end.

Front 54

Pi Bond

Back 54

1. Bond that forms between two p orbitals that overlap side-to-side.

Front 55

A double bond consists of a _________ bond and a ________ bond.
A triple bond consists of a _________ bond and two ______ bonds.

Back 55

1. Sigma
2. Pi

Front 56

Linear VSEPR = ________ hybrid

Back 56

sp

Front 57

Trigonal planar = _______ hybrid

Back 57

sp2

Front 58

Tetrahedral = _______ hybrid

Back 58

sp3

Front 59

Trigonal Bipyramidal = _______ hybrid

Back 59

1. sp3d

Front 60

Octahedral = ________ hybrid

Back 60

1. sp3d2

Front 61

Bond Order using Molecular Orbital

Back 61

1. (Number of electrons in bonding MO) - (number of electrons in antibonding MO)/ 2

Front 62

Molecular Orbital Theory

Back 62

1. Molecular orbitals are simply linear combinations of atomic orbitals.

Front 63

What makes something paramagnetic or diamagnetic?

Back 63

1. Paramagnetic: has unpaired electron
2. Diamagnetic: all electrons are paired.

Front 64

Antibonding orbital

Back 64

1. Molecular orbital that is higher in energy than any of the atomic orbitals from which it was formed.

Front 65

Bonding orbital

Back 65

1. Molecular orbital that is lower in energy than any of the atomic orbitals from which it was formed.

Front 66

A duet is represented by _______ dots.

Back 66

1. Two

Front 67

Octet Rule

Back 67

1. Stable configurations are usually eight electrons in the outermost shell.

Front 68

Resonance Hybrid

Back 68

1. The actual structure of molecule is the intermediate between two resonance structures.

Front 69

Bonds may be viewed as a ______________.

Back 69

1. Continuum

Front 70

Why do atoms bond?

Back 70

1. Maximize attraction force and minimize repulsion force!!
2. To achieve full shell look
3. Obtain lower energy.