Chapter 1: Introduction And Review

Front 1

Isotopes

Back 1

- Atoms with the same number of protons but different numbers of neutrons

Front 2

Mass Number

Back 2

- The sum of protons and neutrons

Front 3

What determines the chemical properties of an atom?

Back 3

- The number of protons in the nucleus and the corresponding number of electrons around the nucleus

Front 4

Electrons show properties of what? Why?

Back 4

- Electrons show properties of both particles and waves because they are small and light.

Front 5

Do electrons in atoms and molecules behave more like particles or waves?

Back 5

-Waves

Front 6

Orbitals

Back 6

- Orbitals are mathematical descriptions that chemists use to explain and predict the properties of atoms and molecules.

- Because of electron density and the Heisenberg uncertainty principle; an orbital is an allowed energy state for an electron, with an associated probability function that defines the distribution of electron density in space.

Front 7

The Heisenberg Uncertainty Principle

Back 7

- We can never determine exactly where the electron is, but we can determine its electron density

Front 8

Electron Density

Back 8

The probability of finding the electron in a particular part of the orbital

Front 9

Where is the electron density the highest?

Back 9

- The nucleus

Front 10

What is higher in energy the 1s or 2s orbital and why?

Back 10

The 2s orbital is higher in energy because most of the 2s electron density is farther from the nucleus than that of the 1s.

Front 11

How many 2p atomic orbitals does the second electron shell have?

Back 11

- Three, one oriented in each of the three spatial directions

- These orbitals are called the 2px, 2py,and 2pz according to their direction along the x, y, or z axis

Front 12

Are the 2p or 2s orbitals higher in energy and why?

Back 12

The 2p orbitals are slightly higher in energy than the 2s, because the average location of the electron in a 2p orbital is farther from the nucleus.

Front 13

Each p orbital consists of how many lobes and where?

Back 13

- Each p orbital consists of two lobes, one on either side of the nucleus, with a nodal plane at the nucleus

Front 14

Nodal Plane

Back 14

- The nodal plane is a flat (planar) region of space, including the nucleus, with zero electron density

Front 15

True or False, the three 2p orbitals have identical energies and why?

Back 15

- True, because they only differ in their spatial orientation

Front 16

Degenerate Orbitals

Back 16

- Orbitals with identical energies

Front 17

The Pauli Exclusion Principle

Back 17

- Each orbital can hold a maximum of two electrons, provided that their spins are paired.

Front 18

How many electrons can the first, second, and third electron shell hold and what orbitals make up each shell?

Back 18

-1st = 2 electrons/ one 1s orbital

- 2nd = 8 electrons/ one 2s orbital and three 2p orbitals

- 3rd = 18 electrons/ one 3s orbital, three 3p orbitals, and five 3d orbitals

Front 19

Hund's Rule

Back 19

- When there are two or more orbitals of the same energy, electrons will go into different orbitals rather than pair up in the same orbital

Front 20

Which elements can have "expanded octets" and why?

Back 20

- Elements in the third and higher rows (such as Al, Si, P, S, Cl, and above)

- Because they have low-lying d orbitals available

Front 21

Nonpolar Covalent Bond

Back 21

- A bond with the electrons shared equally between the two atoms

Front 22

Polar Covalent Bond

Back 22

- An unequally shared pair of bonding electrons

Front 23

What measures bond polarity?

Back 23

- It's dipole moment

Front 24

Dipole Moment

Back 24

- The amount of charge separation multiplied by the bond length

Front 25

Electronegativities

Back 25

- Used to predict if a bond will be polar and the direction of its dipole moment

Front 26

How to calculate formal charges

Back 26

- Count how many electrons contribute to the charge of each atom and compare that number with the number of valence electrons in the free neutral atom.

Front 27

The electrons that contribute to an atom's charge are

Back 27

- All its unshared (nonbonding) electrons; plus

- Half the (bonding) electrons it shares with other atoms, or one electron of each bonding pair.

Front 28

Resonance Structures

Back 28

- Different ways of drawing the same compound

Front 29

What does it mean when a possitive charge is delocalized?

Back 29

- The positive is spread out over both bonding atoms

Front 30

Resonance-stabilized cation

Back 30

- Spreading the positive charge over two atoms makes the ion more stable than it would be if the entire charge were localized only on one atom.

Front 31

If each of the carbon-oxygen bonds are halfway between a single bond and a double bond, they are said to have a bond order of ____?

Back 31

- 1.5

Front 32

Major Resonance Contributor

Back 32

- The more stable resonance form

Front 33

Minor Resonance Contributor

Back 33

- The less stable resonance form

Front 34

Electrostatic Potential Map (EPM)

Back 34

- The electron rich region is red

- The electron poor region is blue

Front 35

What is the goal when drawing resonance forms? What are the best candidates for drawing resonance forms?

Back 35

- To draw structures that are as low in energy as possible

- The best candidates are those that have the maximum number of octets and the maximum number of bonds with a minimum amount of charge separation

Front 36

Only ____ can be delocalized

Back 36

Electrons

Front 37

What is the first rule when drawling resonance structures?

Back 37

- All the resonance structures must be valid Lewis structures for the compound

Front 38

What is the second rule when drawling resonance structures?

Back 38

- Only he placement of the electrons may be shifted from one structure to another.

- Electrons in double bonds and lone pairs are the ones that are most commonly shifted

- Nuclei cannot be moved, and the bond angles must remain the same.

Front 39

What is the third rule when drawling resonance structures?

Back 39

- The number of unpaid electrons (if any) must remain the same

- Most stable compounds have no unpaired electrons, and all the electrons must remain paired in all the resonance forms

Front 40

What is the fourth rule when drawling resonance structures?

Back 40

- The Major resonance contributor is the one with the lowest energy.s

- Good contributors generally have all octets satisfies, as many bonds as possible, and as little charge separation as possible.

- Negative charges are more stable on more electronegative atoms, such as O, N, and S.

Front 41

What is the fifth rule when drawling resonance structures?

Back 41

- Resonance stabilization is most important when it serves to delocalize a charge over two or more atoms

Front 42

Resonance forms can be compared using the following criteria, beginning with the most important:

Back 42

1. As many octets as possible

2. As many bands as possible

3. Any negative charges on electronegative atoms

4. As little charge separation as possible