3B1 Bonding

Question 1

Define:

electronegativity

Answer 1

The attraction of atoms to electrons.

More electronegative atoms pull shared electrons closer, creating dipoles.

Question 2

List the two main types of covalent bonds based on electronegativity.

Answer 2

1. Polar covalent bonds
2. Non-polar covalent bonds

Polar bonds have uneven electron sharing, while non-polar bonds share electrons evenly.

Question 3

List the three types of covalent bonds based on the number of electron pairs shared.

Answer 3

1. Single covalent bond
2. Double covalent bond
3. Triple covalent bond

Single bonds share one electron pair, double bonds share two electron pairs, and triple bonds share three electron pairs.

Question 4

List the two categories of electrons in an atom.

Answer 4

1. Core electrons
2. Valence electrons

Valence electrons are those in the outermost energy shell.

Core electrons are not involved in bonding, while valence electrons are used in chemical bonding.

Question 5

Define:

ionic bonding

Answer 5

Bonding that occurs when one atom donates electrons to another, creating oppositely charged ions that attract each other.

Ionic bonds involve complete transfer of electrons.

Question 6

What is the role of metals in the formation of ionic compounds?

Answer 6

They donate electrons to form cations.

Metals have low electronegativity and prefer to lose electrons.

Question 7

What do nonmetals do during the formation of ionic compounds?

Answer 7

They receive electrons to form anions.

Nonmetals typically have high electronegativity and prefer to gain electrons.

Question 8

What structure do ionic compounds form?

Answer 8

Crystal lattice structure.

This structure is a rigid and has ordered arrangement of ions. An example includes sodium chloride (NaCl).

Question 9

Define:

lattice energy

Answer 9

Energy released when oppositely charged ions form a crystal lattice.

It represents the energy required to separate one mole of ions.

Question 10

How does the lattice energy of an ionic compound affect its properties?

Answer 10

Higher lattice energy increases:

• Melting point
• Boiling point
• Hardness

Lattice energy is proportional to the charge and inversely proportional to the ionic radius.

Question 11

What happens when a solid ionic compound is placed in water?

Answer 11

It dissolves into ions.

Water’s polarity helps to separate the ions, forming an aqueous solution.

Question 12

Fill in the blank:

Ionic compounds when melted or in solution are good _______.

Answer 12

Conductors of electricity.

The mobility of ions in solution or molten form allows for electrical conductivity.

Question 13

What is the effect of ionic charge on the structure of ionic compounds?

Answer 13

Holds ions in a rigid arrangement.

Positive ions align with negative ions in a strict fashion.

Question 14

Why are ionic bonds strong?

Answer 14

Due to the electrostatic force between oppositely charged ions.

This force is proportional to the charge and inversely proportional to the distance between ions.

Question 15

How does the charge of ions affect ionic bond strength?

Answer 15

Higher charges result in stronger ionic bonds.

For instance, MgO (Mg²⁺ and O²⁻) has stronger bonds than NaCl (Na⁺ and Cl⁻).

Question 16

List three properties of ionic compounds.

Answer 16

1. High melting and boiling points.
2. Soluble in water.
3. Good conductors in molten or aqueous states.

High melting and boiling points arise from strong electrostatic forces between ions.

Question 17

Why are ionic compounds brittle?

Answer 17

The rigid lattice structure causes the compound to shatter when force misaligns the ions.

Misalignment brings like charges close together, resulting in repulsion and breakage.

Question 18

True or False:

Ionic compounds are generally polar.

Answer 18

True

The large electronegativity difference between ions results in polarity.

Question 19

What is covalent bonding?

Answer 19

A chemical bond formed by the sharing of electrons between atoms.

Covalent bonds typically form between nonmetals.

Question 20

Why are covalent bonds directional?

Answer 20

Because shared electrons are localized between two specific atoms.

This contrasts with ionic or metallic bonds, which are non-directional.

Question 21

List five differences between ionic and covalent bonds.

Answer 21

1. Ionic bonds transfer electrons; covalent bonds share electrons.
2. Ionic bonds are non-directional; covalent bonds are directional.
3. Ionic compounds are usually solids; covalent compounds can be gases, liquids, or solids.
4. Covalent compounds are molecules; ionic compounds are lattices.
5. Ionic bonds are stronger on average.

Question 22

List three factors that affect bond strength.

Answer 22

1. Bond type (single, double, or triple).
2. Bond length.
3. Electronegativity difference.

Bond strength increases with shorter bond lengths and higher electronegativity differences.

Question 23

How does bond length relate to bond strength in covalent bonds?

Answer 23

Shorter bond lengths generally correspond to stronger bonds.

For example, a triple bond is shorter and stronger than a single bond.

Question 24

How does hybridization affect bond lengths in covalent molecules?

Answer 24

It alters the bond lengths by mixing orbitals.

Single bonds (sp³) are longer and weaker than double bonds (sp²) or triple bonds (sp).

Question 25

What is the typical **bond length** of a **single covalent bond**?

Answer 25

0.15 nm

Question 26

What is the **strongest** type of covalent bond?

Answer 26

A **triple bond**, as it involves three shared electron pairs. ## Footnote Triple bonds, like in N≡N, are shorter and stronger than single or double bonds.

Question 27

Why are triple covalent bonds **shorter than single bonds**?

Answer 27

Triple bonds involve **more electron sharing**, pulling the nuclei closer together. ## Footnote Increased electron density between nuclei enhances attraction and shortens the bond.

Question 28

Why are covalent compounds **poor conductors of electricity**?

Answer 28

Covalent compounds **lack free-moving charged particles**. ## Footnote In contrast, ionic and metallic bonds allow free movement of ions or electrons.

Question 29

# Define: metallic bonding

Answer 29

Bonding that occurs when **metal atoms share a '*sea*' of free-moving electrons**. ## Footnote Metallic bonds only occur between metal atoms.

Question 30

What **percentage of the periodic table** do metals make up?

Answer 30

About **80%**. ## Footnote Metallic characteristics are highest on the left side of the periodic table.

Question 31

What is the **sea of electrons** model?

Answer 31

A model describing the **ability of electrons to move freely** among bonded metal atoms. ## Footnote This movement is crucial for the properties of metals such as electrical conductivity and thermal energy transfer.

Question 32

How does the electron **sea model explain metallic bonding**?

Answer 32

Metal nuclei surrounded by freely moving delocalized electrons.

Question 33

Why are metallic bonds **flexible**?

Answer 33

Because the sea of electrons allows metal atoms to **move without breaking the bond**. ## Footnote This flexibility makes metals malleable and ductile.

Question 34

What is the **crystalline lattice configuration**?

Answer 34

The arrangement of metal atoms in a **structured pattern** for stability. ## Footnote This structure provides stability for electron movement.

Question 35

What does it mean for metals to be **malleable**?

Answer 35

They can be **molded and shaped** for various uses. ## Footnote This property is essential in construction and transportation.

Question 36

# Define: ductility

Answer 36

The ability of a metal to be **stretched into thin wires**. ## Footnote This is important for electrical applications.

Question 37

# Fill in the blank: The **smaller the metal atom**, the \_\_\_\_\_\_ the metallic bonds will be.

Answer 37

stronger ## Footnote Smaller atoms have a stronger attraction between protons and electrons.

Question 38

What happens to metallic characteristics from **left to right on the periodic table**?

Answer 38

They tend to **decrease**. ## Footnote This trend is observed as elements transition to metalloids.

Question 39

# True or False: **Metallic bonds are weaker** than ionic bonds.

Answer 39

False ## Footnote The strength of metallic bonds depends on the number of delocalized electrons and the charge of metal cations.

Question 40

How does metallic bonding **compare** to ionic bonding in **conductivity**?

Answer 40

Metallic bonding allows free electron flow, making **metals good conductors in both liquid and solid states**, unlike ionic solids. ## Footnote Ionic compounds conduct electricity only when molten or dissolved in water.

Question 41

Why are **metals malleable** while ionic compounds are not?

Answer 41

* Metals: They have a **sea of electrons allowing atoms to slide** past each other. * Ionic compounds: They form **rigid lattices that break** under force. ## Footnote The delocalized nature of metallic bonds enables flexibility.

Question 42

What is the relationship between **metallic bonding and alloy formation**?

Answer 42

Metallic bonding allows the **mixture of different metals** to form alloys with enhanced properties. ## Footnote Alloys, such as steel (Fe + C), combine the strength of metals with added flexibility or resistance.