Chemistry Chapter 10- Part 1

Front 1

Describe VSEPR.

Back 1

Valence Shell Electron Pair Repulsion Theory- It is a proposed theory that suggests that the "geometric" arrangement of terminal atoms or groups about central atom is determined by repulsion of electron pairs. Electron groups repel due to Coulombic Forces

Front 2

What determines the geometry of a molecule?

Back 2

-repulsion between electron groups on interior atoms of a molecule

-overlapping orbitals

-# of bonding and nonbonding electron pairs (electron domains) about the central atom

Front 3

Why is it harder to visualize molecules with 4 or more electron groups rather than 2 or 3?

Back 3

*Think about balloons. If you blow up 3 balloons and tie them together, you can lay them on a flat surface, but if there is a 4th balloon, you have to either tie it to the front or the back. This is why adding a 4th electron group makes it more difficult to visualize the molecule. You have to rotate the entire molecule to see all the electron groups.

Front 4

What is Electron Geometry?

Back 4

geometrical arrangement of the electron groups

Front 5

What is Molecular Geometry?

Back 5

3D geometrical arrangement of atoms of a molecule in space based on both bonding and nonbonding electron pairs

Front 6

Name 5 things that lone pairs do.

Back 6

1. They are greater than bonding pairs.

2. They exert a greater repulsion than bonding pairs.

3. They occupy and angular space.

4. They decrease bond angles.

5. They compress bonding pairs.

Front 7

Why is a lone pair more spread out in space?

Back 7

because a lone pair is attracted to one atom's nucleus while bonding electron pairs are attracted to two

Front 8

Explain the AXN method for figuring out geometry of atoms and molecules.

Back 8

When trying to figure out the geometry of a molecule, you can use the AXN method. The 'A' stands for the central atom ( least electronegative atom). The 'X' stands for the number of atoms bonded to the central atom. The 'N' stands for the number of lone pairs.

Front 9

Linear

Back 9

-Linear AX₂ | 180° | nonpolar | sp

Front 10

Trigonal Planar

Back 10

-Trigonal Planar AX₃ | 120° | nonpolar |sp2

-Trigonal Planar Bent AX₂N₁ | <120° | polar| sp2

Front 11

Tetrahedral

Back 11

-Tetrahedral AX₄ | 109° | nonpolar | sp3

-Pyramidal AX₃N₁ | <109.5° | polar |sp3

-Bent AX₂N₂ | <109.5° | polar |sp3

Front 12

Trigonal Bipyramidal

Back 12

-Trigonal Bipyramidal AX₅ | 90°,120°,180° | nonpolar

-Seesaw AX₄N₁ | 90°,120°,180° | polar

-T-shaped AX₃N₂ | 90° & 180° | polar

-Linear AX₂N₃ | 180° | nonpolar

Front 13

Octahedral

Back 13

Octahedral AX₆ | 90° & 180° | nonpolar

Square Pyramidal AX₅N₁ | 90° & 180° | polar

Square Planar AX₄N₂ | 90° & 180° | nonpolar

Front 14

Does the presence of polar bonds mean that the molecule that compose the polar bonds is polar?

Back 14

Absolutely not.

Front 15

True or False?

Formation of Chemical Bond: energy lowered

No Chemical Bond: energy raised

Back 15

True!

Front 16

Describe equillibrium state of electron interaction.

Back 16

This is where there is a minimum amount of interaction energy. The two atomic 1s orbitals have a significant amount of overlap and electrons spend time in the internuclear region so electrons interact with both nuclei.

Front 17

Describe Interaction Energy of Electrons.

Back 17

The interaction energy of electrons is usually negative (or stabilizing) when the interacting atomic orbitals contain a total of 2 electrons that can spin-pair.

Front 18

Describe Electron Overlapping.

Back 18

Two half-filled electron orbitals approach each other and (1)overlap, (2)adapt an opposite spin for themselves, and (3)create a net energy stabilization. This forms a covalent bond.

*Full orbitals are not involved in bonding.

Front 19

Describe Hybrid Orbitals.

Back 19

They are simply a blend of atomic orbitals.

Standard atomic orbitals combine to form new atomic orbitals that correspond more closely to the actual distribution of electrons when atoms bond chemically.

*requires lots of energy to hybridize

*more bonds=more hybrids

Front 20

Describe Electron Probability Density.

Back 20

Electron Probability Density is more concentrated in a single directional lobe allowing the maximum overlaps which minimizes molecular energy.

Front 21

Describe Charge Distribution.

Back 21

Charge Distribution of a molecule is based on the concept of electronegativity difference between atoms making up the bond.

Front 22

Describe Vectors.

Back 22

The vector (direction & magnitude) direction of the dipole moment ( μ )determines bond polarity in a molecule.

Front 23

Describe Valence Bond Theory.

Back 23

When atoms chemically bond, "constructive interference" results from electron sharing of atomic orbitals.

Front 24

What is a sigma (σ ) bond?

Back 24

It is a covalent bond that results when atomic orbitals overlap horizontally to the internuclear axis. Sigma bonds form single bonds.

Front 25

What is a pi ( π ) bond?

Back 25

It is a covalent bond that results when P atomic orbitals overlap perpendicular to the internuclear axis. Pi bonds help form double and triple bonds AND top & bottom atom bonds.

Front 26

If asked to state orbital hybridization of an atoms, how do you explain it?

Back 26

Maksure to state these things in your answer.

- electron domain geometry

- bonds--> sigma and pi? ( sigma bonds can make single bonds and pi bonds can make double or triple bonds)

-What is the atom doing?

example: sp³ C₂H₆ The C is taking the 2s orbital and mixing it with three 2p orbitals which creates four new molecular orbitals called sp³.

Front 27

C-H Bond

Back 27

When hydrogen and carbon is formed, the hydrogen atomic orbital overlaps with the 2pZ carbon atomic orbital along an axis. This process forms a sigma bond.

Front 28

Explain the Dipole.

Back 28

A dipole arises when two electrical charges of equal magnitude but opposite sign are separated by distance.

*For a polyatomic molecule, we treat the dipoles as 3D vectors. The sum of the vectors will give us the dipole for the molecule.

The Dipole Moment (μ ) =

μ = QrWhere Q is the magnitude of the charges and r is the distance.

Front 29

How to determine orbital hybridization of an atom.

Back 29

First, know that sigma bonds are always the 1st bonds and pi bonds are always the second and third bonds.

*However many pi bonds there are, subtract that number from 3.

example: If the element carbon has *1* pi bond in C₂H₄ , the hybridization of the atom is sp² ( 3 - *1*= 2 )

Front 30

Which is weaker? Pi bond or sigma bond?

Back 30

Pi bond is weaker because sigma bonds have more overlap.

Front 31

Which elements can undergo hybridization to form single,double,and triple bonds?

Back 31

C,O,N,P,S

Front 32

How to draw a sigma and pi bond coming together?

Back 32

Front 33

How to draw a pi bond coming together?

Back 33