3. Bonding and Chemical Interactions

Question 1

What are the two types of chemical bonds?

Answer 1

1. Ionic bond (strongest)
2. Covalent Bond

Question 2

What is the octet rule?

Answer 2

Atoms will bond to try to create a stable octet with 8 valence electrons.

Question 3

What are the elements that defy the octet rule, creating an incomplete octet with with fewer than 8 electrons?

Answer 3

H- Hydrogen
He- Helium (2 electrons)
Li-Lithium (2 electrons)
Be- Beryllium (4 electrons)
B- Boron (6 electrons)

Question 4

What are the elements that defy the octet rule, creating an expanded number of electrons?

Answer 4

All elements in period three, which can extend their shell to include d orbitals.

Essentially, transitional metals

Question 5

What are elements that also defy the octet rule because of the number of electrons?

Answer 5

Those with an odd number of electrons.
For example, NO (nitric oxide) has 11 valence electrons

Question 6

How are ionic bonds and covalent bonds different?

Answer 6

Ionic: From between ions and involve the gain or loss of electrons
Covalent: occur when electrons are shared between atoms

Question 7

Which periodic trend determines whether a covalent bond is polar or non-polar?

Answer 7

The polarity in a covalent bond is determined by differences in electronegativity between the two atoms.

Question 8

How do you define an ionic bond?

Answer 8

formed via the transfer of one or more electrons from an element with a relatively low ionization energy to an element with a relatively high electron affinity

There is electron transfer, NOT sharing

Question 9

What is the difference in electronegativity between two elements that result in an ionic bond?

Answer 9

Ionic bonds occur between elements with large differences in electronegativity (∆EN>1.7), usually between metals and nonmetals.

Question 10

What are crystalline lattices? What type of bond (covalent or ionic) forms them and why?

Answer 10

They are large, organized arrays of ions formed by ionic compounds

Why? The attractive forces between oppositely charged ions are maximized, and the repulsive forces between the ions of like charge are minimized.

Question 11

What are unique properties of ionic bonds?

Answer 11

1. Ionic compounds tend to dissociate in water and other polar solvents
2. Ionic solids tend to have high melting points
3. Good conductors of heat and electricity
4. Create crystal lattice arrangements to minimize repulsive forces
5. Large electronegative differences

Question 12

How do you define a covalent bond?

Answer 12

They form via the sharing of electrons between two elements of similar electronegativities

Question 13

Why create a covalent bond rather than an ionic bond?

Answer 13

When two atoms of similar tendency to attract electrons form a compound, it is energetically unfavorable to create ions, so they share.

Question 14

What does the bond order refer to?

Answer 14

Refers to whether a covalent bond is a single bond, double bond, or triple bond?

Question 15

What are the trends as covalent bonds create more bonds (i.e. single, then double, than triple)?

Answer 15

As you create more bonds:
1. The bond length decreases, they get closer
2. They strengthen the bond, i.e. it takes more energy to break the bond

Question 16

What is polarity and in what type of bond does it occur?

Answer 16

It occurs in covalent bond.
It occurs when two atoms have enough electronegative difference to draw electrons closer to one element as opposed to another, but not enough to create an ionic bond.

Question 17

What’s the difference in electronegativities that polar covalent bonds occur?

Answer 17

∆EN= .5-1.7

Question 18

When will covalent bonds be non-polar?
What are some examples?

Answer 18

They will occur when the ∆EN<0.5, or when the elements have the exact same electronegativity. A lot of these occur in group 7…

H2, N2, O2, F2, Cl2, Br2, I2

Question 19

What is a dipole?

Answer 19

This is created in polar covalent bonds when one atom gets more of the shared electron than the other.

Question 20

What is the equation for dipole moment and what does it mean?

Answer 20

p=qd

p=dipole moment
q= magnitude of the charge
d= displacement vector separating the two particle charges.

Measured in Debye Units (coulomb-meters)

It is a vector quantity of a dipole

Question 21

What are coordinate covalent bonds?

Answer 21

They result when a single atom provides both bonding electrons, while the other atom does not contribute any.
This commonly occurs in equations where on atom is connected to H+ (i.e. the H+ atom has NO electrons, so the other atom has to give both electrons to them)

These are also called nucleophile-electrophile reactions.

Question 22

What are the steps to draw a Lewis structure?

Answer 22

1. Draw out the backbone of the structure, the least electronegative element is the middle element. The Halogens (F, Cl, Br, and I) always occupy a terminal position
2. Count the number of valence electrons and add them together.
3. Draw a single bond between the central atom and the atoms surrounding it (each bond counts as 2 electrons).
4. Complete the octets of all atoms bonded to the central atom (except the central atom), using the remaining valence electrons
5. Put the remaining atoms on the central atoms. If there are no valence electrons left, form bonds with the other atoms paired electrons (coordinate covalent bonds)

Question 23

What is a formal charge, and what is the equation?

Answer 23

This is done to determine if a Lewis Structure is representative of the actual arrangement of atoms in a compound. They exist when an atom is surrounded by more or fewer valence electrons that it has in its neutral state (assuming equal sharing of electrons in a bond

The charge of the ion or compound is equal to the sum of the formal charges of the individual atoms comprising the ion or compound.

Question 24

What is the formal charge equation, and what’s the cheat way?

Answer 24

Formal Charge= number of valence electrons given by column- dots - sticks

*** The charge of the ion or compound is equal to the sum of the formal charges of the individual atoms comprising the ion or compound.

Question 25

What are resonance structures?

Answer 25

They are alternate arrangements of Lewis structures that demonstrate the same arrangement of atoms, but differ in the specific placement of electrons. A more complicated definition: include any molecule with a π (pi) system of electrons; These represent all of the possible configurations of electrons- stable and unstable, that contribute to the overall structure.

Question 26

What are some of the guidelines that are used to assess the stability (i.e. identify the resonance structure with the highest likelihood of occurring)?

Answer 26

1. A Lewis structure with small or no formal charge is preferred over a Lewis structure with a large formal charge 2. A Lewis structure with less separation between opposite charges is preferred over a Lewis structure with a large separation of opposite charges 3. A Lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the negative formal charges are placed on less electronegative atoms.

Question 27

What is the Valence Shell Electron Pair Repulsion Theory?

Answer 27

predicts the three-dimensional molecular geometry of covalently bonded molecules. In this theory, electrons- whether bonding or non-bonding- arrange themselves to be as far apart as possible from each other in three-dimensional space, leading to characteristic geometries. **This is different from resonance structures**

Question 28

How do you determine the structure (3-d, non-lewis) of molecules?

Answer 28

1. Draw the Lewis Structure 2. Count the total number of electrons and non-binding electron pairs in the valence shell of the central atom. 3. Arrange the electron pairs around the central atom so that they are as far apart as possible. For example, the compound AX2, has the Lewis structure X:A:X. To position these electron pairs as far apart as possible, their geometric structure should be linear. 4. You also must take note of the electron pairs of the central atom. If there are electron pairs, treat them like another atom to create the shape (just be cognizant to not add in the final diagram)

Question 29

What's the difference between **electronic geometry** and **molecular geometry**?

Answer 29

**Electronic Geometry**: refers to the position of all *electrons* in molecule, whether bonding or nonbonding. **Molecular geometry**: refers to the position of *only the bonding pairs of electrons* in a molecule

Question 30

What is the definition of a **molecular orbital**?

Answer 30

This describes the probability of finding the bonding electron in a given space

Question 31

What are sigma and Pi bonds in a covalent bond?

Answer 31

describe the patterns of overlap observed when molecular bonds are formed Sigma (σ) Bonds: Result of head to head overlap of two orbitals Pi (π) Bonds: the result of the overlap of two parallel electron cloud densities Layman's Terms: Sigma bond: 1 sigma bond = single, double, or triple bond between two atoms. There is only one sigma bond, even if it is a triple bond. Pi bond: 1-sigma bond. In a single bond, there is no pi bond. In a double bond there is one pi bond. In a triple bond, there are 2 pi bonds.

Question 32

If two atoms form a triple bond, how many pi bonds are there? How many sigma bonds are there?

Answer 32

**Sigma (σ) Bonds**: 1 sigma **Pi (π) Bonds**: 2 pi bonds. Layman's Terms: Sigma bond: 1 sigma bond = single, double, or triple bond between two atoms. There is only one sigma bond, even if it is a triple bond. Pi bond: 1-sigma bond. In a single bond, there is no pi bond. In a double bond there is one pi bond. In a triple bond, there are 2 pi bonds.

Question 33

What are some characteristics of covalent bonds?

Answer 33

1. Lower melting and boiling points compared to ionic compounds. 2. Poor conductors of electricity in liquid or aqueous state.

Question 34

Describe the relationship between bond strength, bond length, and bond energy.

Answer 34

Bond strength is defined by the electrostatic attraction between nuclei and electrons; multiple bonds (higher bond order) increases strength. Bond length is a consequence of these attractions. The stronger the bond, the shorter it is. Bond energy is the minimum amount of energy needed to break a bond. The stronger the bond, the higher the bond energy.

Question 35

For what values of **∆EN** will a non-polar covalent bond form? Polar Covalent? Ionic? I.e. What's the DIFFERENCE in electronegativity for these bonds to form

Answer 35

**Non-Polar Covalent**: ∆EN < .5 **Polar Covalent**: ∆EN .5-1.7 **Ionic**: 1.7 and above

Question 36

Draw a Lewis dot structure for the carbonate ion (CO₃^-2) and its two other resonance structures.

Answer 36

1. Draw out the backbone of the structure, the least electronegative element is the middle element. The Halogens (F, Cl, Br, and I) always occupy a terminal position 2. Count the number of valence electrons and add them together. 3. Draw a single bond between the central atom and the atoms surrounding it (each bond counts as 2 electrons). 4. Complete the octets of all atoms bonded to the central atom (except the central atom), using the remaining valence electrons 5. Put the remaining atoms on the central atoms. If there are no valence electrons left, form bonds with the other atoms paired electrons (coordinate covalent bonds)

Question 37

Predict the molecular geometries of the following molecules. PCl₅ MgF₂ AlF₃ UBr₆ SiH₄

Answer 37

PCl₅- Triangle pyramid MgF₂- Linear AlF₃ – Triagonal Planar UBr₆- Octahedral SiH₄ -Tetrahedral

Question 38

What is the shape between x-a-x?

Answer 38

Linear

Question 39

What is the shape between AB₃?

Answer 39

**Trigonal Planar**, angle 120 degrees

Question 40

What is the shape between AB₄?

Answer 40

**Tetrahedral** Angle 109.5

Question 41

What is the shape between AB₅?

Answer 41

**Trigonal bipyramidal**, angles 90, 120, 180

Question 42

What is the shape between AB₆?

Answer 42

**Octahedral**, angles 90, 180

Question 43

If you have two compounds, NaCl and CO₂, which would have the higher boiling point?

Answer 43

NaCl because it is an **ionic compound**. They have much higher boiling points than ionic compounds. It takes more energy to break an ionic bond than a covalent bond.

Question 44

When using the VSEPR Theory, what's important to keep in mind? (i.e. electron pairs)

Answer 44

Keep in mind that paired electrons that do not form bonds also need to be accounted for. Even though they won't make a physical part of the diagram, it will influence the shape of the final diagram

Question 45

What are three types of intermolecular forces?

Answer 45

London Dispersion Forces, dipole-dipole interactions, and hydrogen bonds.

Question 46

What are London Dispersion Forces? How strong are they? What can influence them?

Answer 46

They are a type of intermolecular force. These are very short lived moments, that occur when electrons are very randomly floating around. In two different molecules, one molecules electron may be in one area, and the other in the other area, creating kind of like a dipole. It’s this that causes a very, very momentary interaction. This also depends on the degree and ease which molecules can be polarized, or how electrons can be shifted around. Large molecules are more easily polarizable than comparable smaller molecules, and thus posses greater dispersions forces. They are the reason why noble gases liquefy at low temperatures.

Question 47

What are dipole-dipole interactions? In what phases do you see them?

Answer 47

They are an **intermolecular force**. Occur between the oppositely charged ends of polar molecules. Stronger than London dispersion forces. Evident in the **solid and liquid phases**, but negligible in the gas phase due to the distance between the particles.

Question 48

What are hydrogen bonds?

Answer 48

They are NOT bonds, but a type of intermolecular force. This is a very strong type of dipole-dipole interactions, and only occur between fluorine, oxygen, and nitrogen Have only about **10% of the strength of a covalent bonds**, so the electrostatic interactions can be overcome with small or moderate amounts of energy Specialized subset of dipole-dipole interactions involved in intra-and intermolecular attraction; **hydrogen bonding** occurs when hydrogen is bonded to one of three very electronegative atoms- **fluorine, oxygen, or nitrogen**. Substances that do have hydrogen bonds tend to have unusually **high boiling points**. DNA base pairs (AGCT) use hydrogen bonds

Question 49

In what three elements to hydrogen bonds occur in?

Answer 49

fluorine, oxygen, or nitrogen.

Question 50

Why are intermolecular forces important?

Answer 50

These bonds are important because they keep a substance together in its **solid or liquid state** and determine whether two substances are **miscible or immiscible in a solution**. Miscible: Two liquids that appear to mix completely together are said to be miscible

Question 51

Rank the major intermolecular forces from strongest to weakest

Answer 51

(Strongest) Hydrogen bond > dipole-dipole interaction > London Dispersion forces (Weakest)

Question 52

Describe what occurs during dipole-dipole interactions.

Answer 52

A dipole consists of a segment of a molecule with partial positive and partial negative reactions. The positive end of one molecule is attracted to the negative end of another molecule, and visa versa. Do not occur in non-dipoles

Question 53

In order to exhibit hydrogen bonding, what must be true of a given molecule?

Answer 53

It must have **fluorine, oxygen or nitrogen** (very electronegative atoms)