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Chemistry Module 1 > Bonding > Flashcards

Bonding Flashcards

What binds atoms together in elements and compounds? (89 cards)

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1
Q

Define electronegativity.

A

Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond.

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2
Q

How does electronegativity affect bond formation?

A

A large difference in electronegativity results in an ionic bond, while a smaller difference creates covalent bonds.

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3
Q

Explain the concept of ionization energy.

A

Ionization energy is the energy required to remove an electron from an atom or ion in the gas phase.

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4
Q

How does ionization energy relate to bond formation?

A

High ionization energy in metals makes it easier to form cations, essential for ionic bonding with non-metals.

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5
Q

Why are covalent bonds directional?

A

Covalent bonds are directional because they involve shared electron pairs between specific atoms, dictating the bond’s orientation.

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6
Q

Explain ionic bonding.

A

Ionic bonding occurs when one atom transfers electrons to another, creating oppositely charged ions that are held together by electrostatic attraction.

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7
Q

Why do metals form cations in ionic bonds?

A

Metals have low ionization energies, making it easier to lose electrons and form positive ions that bond with negative ions.

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8
Q

Why do non-metals form anions in ionic bonds?

A

Non-metals have high electronegativity, making it easier for them to gain electrons and form negatively charged ions.

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9
Q

Define the concept of a polar covalent bond.

A

A polar covalent bond is formed when electrons are shared unequally between two atoms due to differing electronegativities.

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10
Q

Why are polar bonds considered partial charge bonds?

A

The electron distribution is uneven, creating a dipole, with one atom having a partial negative charge and the other a partial positive charge.

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11
Q

Explain the significance of a non-polar covalent bond.

A

In a non-polar covalent bond, electrons are shared equally, typically occurring between identical atoms with equal electronegativity.

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12
Q

What is the concept of bond polarity?

A

Bond polarity refers to the unequal sharing of electrons between two atoms, causing a dipole moment in a molecule.

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13
Q

How does the position of elements in the periodic table influence bond polarity?

A

Elements on the left side (metals) have lower electronegativity, while elements on the right (non-metals) have higher electronegativity, influencing bond polarity.

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14
Q

How do you determine the type of bond based on electronegativity difference?

A

A difference greater than 1.7 typically forms ionic bonds, between 0.4-1.7 forms polar covalent bonds, and less than 0.4 forms non-polar covalent bonds.

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15
Q

Explain what is meant by the term ‘bond energy’.

A

Bond energy is the amount of energy required to break a bond in one mole of a molecule, indicating the strength of the bond.

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16
Q

How does bond length relate to bond strength?

A

Shorter bonds generally have higher bond energy, as the atoms are closer and more tightly held together.

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17
Q

Define the term ‘valency’.

A

Valency is the number of electrons an atom needs to gain, lose, or share in order to achieve a stable electron configuration.

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18
Q

Why is valency important in determining molecular formulas?

A

Valency helps in determining how atoms combine in correct proportions to form molecules with stable electron configurations.

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19
Q

Describe the role of Lewis dot diagrams in understanding bonding.

A

Lewis dot diagrams show the valence electrons of atoms, helping visualize how atoms bond by sharing or transferring electrons.

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20
Q

What is the octet rule?

A

The octet rule states that atoms tend to form bonds in order to achieve a stable electron configuration with eight electrons in their valence shell.

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21
Q

Why do atoms form bonds?

A

Atoms form bonds to achieve a more stable electron configuration, typically by completing their outer electron shell.

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22
Q

What is meant by the term ‘ionic lattice’?

A

An ionic lattice is a three-dimensional arrangement of ions in a regular, repeating pattern held together by electrostatic forces.

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23
Q

How do ionic bonds affect the physical properties of ionic compounds?

A

Ionic bonds create strong electrostatic forces that result in high melting and boiling points, and the ability to conduct electricity when dissolved in water.

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24
Q

What is a covalent network solid?

A

A covalent network solid is a structure where atoms are bonded by a continuous network of covalent bonds (e.g., diamond or silicon dioxide).

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25
How does the structure of covalent network solids influence their properties?
The strong covalent bonds throughout the structure result in high melting points and hardness, as well as electrical non-conductivity.
26
What is the difference between covalent molecular compounds and covalent network solids?
Covalent molecular compounds consist of discrete molecules held together by intermolecular forces, while covalent network solids have atoms bonded in an extensive network throughout the material.
27
Define metallic bonding.
Metallic bonding occurs when metal atoms release some of their electrons, creating a 'sea' of delocalized electrons that move freely between positive metal ions.
28
What is the importance of delocalized electrons in metallic bonding?
Delocalized electrons in metallic bonding allow metals to conduct electricity and heat, and contribute to their malleability and ductility.
29
How does metallic bonding explain the malleability and ductility of metals?
The delocalized electrons allow metal ions to slide past each other without breaking bonds, making metals flexible and capable of being drawn into wires.
30
What is an example of a covalent molecular substance?
Water (H₂O) is a covalent molecular substance, where molecules are held together by hydrogen bonds.
31
Why are ionic compounds typically solid at room temperature?
The strong electrostatic forces between oppositely charged ions in an ionic lattice result in a solid structure at room temperature.
32
What is a characteristic feature of molecular compounds?
Molecular compounds typically have lower melting and boiling points than ionic compounds because their intermolecular forces are weaker.
33
Define the concept of hydrogen bonding.
Hydrogen bonding is a strong type of dipole-dipole interaction occurring when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine.
34
Why does water have a high boiling point?
The hydrogen bonds between water molecules are strong and require significant energy to break, resulting in a high boiling point.
35
What is the difference between intermolecular and intramolecular forces?
Intermolecular forces act between molecules, while intramolecular forces hold atoms together within a molecule (e.g., ionic or covalent bonds).
36
Why do ionic compounds have high melting points?
The strong electrostatic attraction between the ions in an ionic lattice requires a large amount of energy to break apart.
37
Explain the concept of an alloy.
An alloy is a mixture of two or more metals, which often have superior properties compared to pure metals.
38
Why are alloys typically stronger than pure metals?
The different sizes of atoms in an alloy prevent the layers of metal ions from sliding past each other easily, making the material stronger.
39
How does the metallic bond explain the electrical conductivity of metals?
Delocalized electrons in metallic bonds move freely through the lattice, allowing metals to conduct electricity.
40
Why are ionic compounds poor conductors in solid form?
In the solid state, ions are fixed in place and cannot move to conduct electricity, though they can in molten form or when dissolved in water.
41
What is lattice energy?
Lattice energy is the energy released when oppositely charged ions come together to form a lattice in an ionic solid.
42
Why is lattice energy higher for smaller ions?
Smaller ions have a higher charge density, leading to stronger electrostatic attraction between them, which results in higher lattice energy.
43
How does the size of ions affect the strength of ionic bonds?
Smaller ions can pack more closely together, leading to stronger ionic bonds due to increased electrostatic attraction.
44
Why do covalent bonds not conduct electricity?
Covalent bonds do not conduct electricity because there are no free-moving charged particles (electrons or ions) to carry a charge.
45
Define the term 'electronegativity difference' in bond formation.
Electronegativity difference refers to the difference in electronegativity values between two bonded atoms, determining the bond's ionic or covalent nature.
46
How does bond polarity affect the solubility of compounds?
Polar bonds are more likely to dissolve in polar solvents, while non-polar bonds dissolve in non-polar solvents due to similar intermolecular forces.
47
Why is the octet rule not always followed?
The octet rule doesn't apply in cases where elements can expand their valence shells (such as in transition metals) or when certain molecules (like H₂) can only achieve 2 electrons.
48
How do you predict bond polarity using the periodic table?
The difference in electronegativity between the two elements in a bond can be predicted using their positions on the periodic table, with metals having low electronegativity and non-metals having high electronegativity.
49
Why does fluorine have the highest electronegativity?
Fluorine is the most electronegative element because it has a small atomic radius and a high effective nuclear charge, attracting electrons strongly.
50
What is a covalent network structure?
A covalent network structure is a solid where atoms are connected in a network by covalent bonds, such as in diamond or graphite.
51
What is the role of intermolecular forces in determining boiling points?
Stronger intermolecular forces (like hydrogen bonds) require more energy to overcome, resulting in higher boiling points.
52
Why does the boiling point of noble gases increase down the group?
As atomic size increases down the group, the number of electrons increases, enhancing London dispersion forces and raising boiling points.
53
How do metallic bonds differ from ionic and covalent bonds?
Metallic bonds involve delocalized electrons between metal atoms, whereas ionic and covalent bonds involve electron transfer or sharing between specific atoms.
54
55
What is an ionic bond?
An ionic bond is formed when one atom donates an electron to another atom, resulting in the formation of oppositely charged ions that attract each other through electrostatic forces. This bond typically forms between metals and non-metals, where the metal becomes a cation (positive) and the non-metal becomes an anion (negative).
56
How are ionic bonds created?
Ionic bonds are created when atoms with large electronegativity differences (such as metals and non-metals) transfer electrons. The metal loses electrons to form a positively charged ion, and the non-metal gains electrons to form a negatively charged ion, which then attract each other through strong electrostatic forces.
57
What are the properties of ionic compounds?
Ionic compounds tend to have high melting and boiling points, conduct electricity when molten or dissolved in water, and are usually soluble in polar solvents like water. They form crystalline structures and are brittle due to the rigid lattice arrangement of ions.
58
What is a covalent bond?
A covalent bond is formed when two atoms share one or more pairs of electrons to achieve a stable electron configuration. This bond typically occurs between non-metal atoms with similar electronegativities.
59
How are covalent bonds created?
Covalent bonds are formed when two atoms share one or more pairs of electrons. The atoms may have similar or identical electronegativity, leading to the equal or unequal sharing of electrons. The bond results in the formation of a molecule.
60
What are the properties of covalent compounds?
Covalent compounds generally have lower melting and boiling points compared to ionic compounds. They do not conduct electricity in any state because there are no free-moving charged particles. They are often soluble in non-polar solvents and may be gases, liquids, or solids.
61
What is a metallic bond?
A metallic bond is formed when metal atoms release some of their electrons, creating a 'sea' of delocalized electrons that are free to move through a lattice of positive metal ions. This bond is characteristic of metallic elements.
62
How are metallic bonds created?
Metallic bonds are created when metal atoms lose their outermost electrons, which become free-moving within a 'sea' of electrons. The remaining positive metal ions are held together by the electrostatic attraction to the delocalized electrons, forming a metallic structure.
63
What are the properties of metallic compounds?
Metallic compounds have high electrical and thermal conductivity, malleability, and ductility due to the delocalized electrons that can move freely throughout the metal structure. They are also usually shiny and have high melting and boiling points.
64
What is a hydrogen bond?
A hydrogen bond is a type of dipole-dipole interaction that occurs when a hydrogen atom, which is covalently bonded to a highly electronegative atom (such as oxygen, nitrogen, or fluorine), interacts with another electronegative atom in a different molecule.
65
How are hydrogen bonds created?
Hydrogen bonds are created when a hydrogen atom that is covalently bonded to an electronegative atom (like oxygen) is attracted to another electronegative atom (such as nitrogen or fluorine) in a different molecule, creating an intermolecular bond.
66
What are the properties of substances with hydrogen bonds?
Substances with hydrogen bonds tend to have high melting and boiling points due to the strong intermolecular forces. Water, for example, has a high boiling point relative to its molecular weight because of hydrogen bonding.
67
What is a dipole-dipole interaction?
Dipole-dipole interactions occur between polar molecules, where the positive end of one molecule is attracted to the negative end of another molecule. These interactions are stronger than London dispersion forces but weaker than hydrogen bonds.
68
How are dipole-dipole interactions created?
Dipole-dipole interactions are created between polar molecules where one end of the molecule has a partial positive charge and the other end has a partial negative charge. The positive end of one molecule attracts the negative end of another molecule.
69
What are the properties of substances with dipole-dipole interactions?
Substances with dipole-dipole interactions have higher boiling and melting points compared to non-polar molecules due to the attractive forces between the positive and negative ends of polar molecules.
70
What is a London dispersion force?
London dispersion forces are weak, temporary forces that arise from the movement of electrons within atoms and molecules, creating instantaneous dipoles. These forces are present in all molecules but are the only intermolecular force in non-polar molecules.
71
How are London dispersion forces created?
London dispersion forces are created when the movement of electrons in an atom or molecule creates a temporary dipole, which induces a temporary dipole in a neighboring molecule. This interaction is weakest but can become significant in large molecules.
72
What are the properties of substances with London dispersion forces?
Substances with only London dispersion forces tend to have low boiling and melting points. The strength of these forces increases with the size and number of electrons in the molecule, making larger molecules exhibit stronger London dispersion forces.
73
How do the strength of intermolecular forces affect the physical properties of substances?
The stronger the intermolecular forces, the higher the boiling and melting points of a substance. Substances with hydrogen bonds or dipole-dipole interactions have higher boiling points than those with only London dispersion forces.
74
What is the relationship between electronegativity and bonding?
Electronegativity differences between atoms determine the type of bond. If the difference is large, an ionic bond is likely to form, whereas a small difference results in a covalent bond. If the atoms are identical, a non-polar covalent bond forms.
75
What is an allotrope?
An allotrope is a different form of the same element in the same physical state, where the atoms are bonded in different ways, leading to different physical and chemical properties.
76
How do allotropes of an element differ?
Allotropes differ in the arrangement and bonding of atoms, which results in distinct physical and chemical properties despite being made from the same element.
77
Give an example of carbon allotropes.
Carbon has several allotropes, including diamond, graphite, graphene, and fullerenes. These allotropes differ in their atomic structures and have vastly different properties.
78
What is the atomic structure of diamond?
In diamond, each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a rigid, three-dimensional network.
79
What are the physical properties of diamond?
Diamond is extremely hard, has a high melting point, is transparent, and is a poor conductor of electricity due to its strong covalent bonds in a rigid network.
80
What is the atomic structure of graphite?
In graphite, each carbon atom is bonded to three others in flat, two-dimensional layers of hexagonal rings, with weak forces between the layers.
81
What are the physical properties of graphite?
Graphite is soft, slippery (due to weak forces between layers), and an excellent conductor of electricity. Its layers can slide over one another, making it useful as a lubricant.
82
What is graphene?
Graphene is a single layer of carbon atoms arranged in a hexagonal lattice. It is a two-dimensional allotrope of carbon with extraordinary electrical, thermal, and mechanical properties.
83
What are the properties of graphene?
Graphene is incredibly strong, lightweight, transparent, and an excellent conductor of electricity and heat. It is often used in advanced electronics and materials science.
84
How do the properties of graphite and diamond differ?
Diamond is hard, transparent, and non-conductive, while graphite is soft, opaque, and conductive. This difference is due to the different arrangements of carbon atoms and types of bonding.
85
How does the structure of diamond influence its properties?
Diamond's three-dimensional covalent network results in its hardness, high melting point, and non-conductivity, as each carbon atom is bonded to four others, creating a rigid structure.
86
How does the structure of graphite influence its properties?
Graphite's planar structure with weak forces between layers allows for its softness, lubricating properties, and conductivity, as electrons can move freely within the layers.
87
What is the relationship between allotropy and chemical bonding?
Allotropes of elements like carbon, sulfur, and oxygen display different bonding patterns (e.g., covalent network, molecular, etc.), which result in different physical and chemical properties.
88
Why does diamond conduct electricity poorly?
Diamond has no free electrons to carry charge, as all electrons are tightly bonded in the covalent network, preventing electrical conductivity.
89
How does the allotrope structure of carbon contribute to the difference in properties between diamond and graphite?
In diamond, the strong covalent bonds in a 3D structure lead to high hardness and insulating properties, while graphite’s 2D layers of weakly bonded carbon atoms allow for conductivity and lubrication.

Chemistry Module 1 (4 decks)

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