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Bonding And Structure Flashcards

1
Q

What happens to metals during ionic bonding?

A

Metals lose electrons to achieve the nearest noble gas electronic configuration, becoming positive ions.

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2
Q

What do non-metals do in ionic bonding?

A

Non-metals gain electrons to achieve the nearest noble gas electronic configuration, becoming negative ions.

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3
Q

What is ionic bonding?

A

Ionic bonding is the transfer of electrons from metals to non-metals, resulting in the formation of oppositely charged ions.

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4
Q

What is the electrostatic force in ionic bonding?

A

There is a strong electrostatic force of attraction between oppositely charged ions.

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5
Q

Provide an example of ionic bonding.

A

Sodium fluoride (NaF) is an example of ionic bonding.

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6
Q

What happens to the electron in sodium fluoride?

A

The electron is transferred from the metal (sodium) to the non-metal (fluoride).

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7
Q

What structure do ions form in ionic bonding?

A

Ions form a giant ionic lattice structure.

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8
Q

What kind of balance exists in a giant ionic lattice?

A

There is a balance of forces due to constant repulsion between ions of the same charge and attraction between ions of opposite charge.

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9
Q

What is the structure of ionic compounds?

A

They have a giant ionic lattice structure

This structure consists of a regular arrangement of ions held together by strong electrostatic forces.

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10
Q

What are the melting and boiling points of ionic compounds like?

A

They have high melting points and boiling points

The melting point value is influenced by the strength of the attractions between ions.

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11
Q

What is required to overcome the electrostatic attractions in ionic compounds?

A

A lot of energy

The energy required is due to the strong attractions between oppositely charged ions.

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12
Q

What factors influence the strength of ionic bonding?

A

Factors include the charge of the ions and the distance between them

Higher charges and shorter distances typically increase ionic bond strength.

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13
Q

Do ionic compounds conduct electricity in the solid state?

A

No, they do not conduct electricity in the solid state

In solid form, ions are in a fixed position and cannot move.

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14
Q

Under what conditions do ionic compounds conduct electricity?

A

They conduct electricity in the molten state or in aqueous solution

In these states, ions are free to move, allowing conductivity.

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15
Q

What happens to ionic compounds when they dissolve in water?

A

They are ‘pulled apart’ by water molecules

The ions make attractions to different atoms in water, leading to dissolution.

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16
Q

What is the difference between NaCl in solid form and NaCl in aqueous solution?

A

NaCl: solid is a fixed lattice, NaCl: aqueous is dissolved in water

In aqueous solution, Na+ and Cl- ions are free to move.

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17
Q

What is a covalent bond?

A

A covalent bond occurs between two non-metal atoms that share a pair of electrons.

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18
Q

What do the atoms involved in a covalent bond share?

A

The atoms share some of their outer (valence) electrons.

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19
Q

What is the charge of the atoms involved in a covalent bond?

A

Neutral, as no electrons are lost or gained.

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20
Q

What do dot and cross diagrams represent?

A

The outer (valence) electrons of the elements as crosses and dots.

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21
Q

In a covalent bond between hydrogen atoms, how many outer electrons does each hydrogen atom have?

A

Each hydrogen atom has a single outer electron.

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22
Q

In hydrogen fluoride, how many outer electrons does hydrogen have?

A

1 outer electron.

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23
Q

In hydrogen fluoride, how many outer electrons does fluorine have?

A

7 outer electrons.

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24
Q

What is the Octet Rule?

A

Atoms tend to achieve a stable configuration with 8 electrons in their outer shell.

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25
In the example of hydrogen fluoride, how many electrons does fluorine have in its outer shell after forming a covalent bond?
8 electrons.
26
True or False: In a covalent bond, electrons are lost or gained.
False.
27
Fill in the blank: A covalent bond occurs between two _______ atoms.
non-metal
28
What is co-ordinate bonding also known as?
Dative covalent bonding.
29
In co-ordinate bonding, what does one atom provide?
Both the electrons needed to form a covalent bond.
30
What must the atom donating the electrons have?
A lone pair of electrons.
31
What is a lone pair of electrons?
Two electrons not used for bonding.
32
What must the atom receiving the electrons have?
A vacant orbital.
33
What does a vacant orbital mean?
Space for 2 electrons in their outer shell.
34
How is a co-ordinate bond represented?
With an arrow instead of a line.
35
Fill in the blank: Co-ordinate bonding is when one atom provides both the electrons needed to form a _______.
covalent bond.
36
True or False: In co-ordinate bonding, both atoms contribute electrons equally.
False.
37
What is the role of the atom with a lone pair in co-ordinate bonding?
To donate both electrons for the bond.
38
What type of bond is formed when one atom provides both electrons?
Co-ordinate bond.
39
What determines the 3D shape a molecule adopts?
The number of bonding pairs of electrons and the number of lone pairs ## Footnote The presence of lone pairs distorts the shape of the molecule.
40
What theory explains the shape adopted by molecules due to electron pairs?
Electron pair repulsion theory (VSEPR) ## Footnote This theory states that electron pairs repel each other, leading to shapes that minimize repulsion.
41
How do lone pairs influence the shape of a molecule?
They influence the shape though they cannot be 'seen' in the final shape ## Footnote Lone pairs cause adaptations in the basic shapes of molecules and affect bond angles.
42
What is the order of repulsion strength among electron pairs?
LP-LP repulsion > BP-LP Repulsion > BP-BP repulsion ## Footnote LP = Lone Pair, BP = Bond Pair.
43
What is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs?
Linear, 180° ## Footnote The arrangement is Y - X - Y.
44
What is the shape and bond angle of a molecule with 3 bonding pairs and 0 lone pairs?
Trigonal Planar, 120° ## Footnote This shape has three atoms bonded to a central atom.
45
What is the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs?
Tetrahedral, 109.5° ## Footnote In this arrangement, four atoms are bonded to the central atom.
46
What is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pair?
Trigonal Pyramid, 107° ## Footnote The lone pair is positioned above the central atom, affecting the bond angle.
47
What is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs?
Non-linear, 104.5° ## Footnote The presence of two lone pairs compresses the angle between the bonded pairs.
48
What is the shape and bond angle of a molecule with 6 bonding pairs and 0 lone pairs?
Octahedral, 90° ## Footnote This shape has six atoms bonded to a central atom.
49
How much does each lone pair reduce the bond angle?
Approximately 2.5° ## Footnote This reduction affects the overall geometry of the molecule.
50
In an F-F bond, how is the electron density distributed?
The electron density is distributed equally/evenly.
51
What is electronegativity?
The ability of an atom to attract a pair of electrons in a covalent bond.
52
What scale is used to grade electronegativity?
The Pauling Scale.
53
What factors determine the electronegativity of an atom? List them.
* The nuclear charge (number of protons) of the atom * The atomic radius of the atom * The number of principal energy levels the atom has.
54
True or False: In a covalent bond, both atoms have the same electronegativity.
False.
55
What happens to the attraction in a covalent bond when the nuclei of the atoms are identical?
The attraction to the electrons is equal.
56
What is the relationship between atomic radius and electronegativity?
The smaller the atomic radius, the greater the electronegativity.
57
What is the relationship between nuclear charge and electronegativity?
The higher the nuclear charge, the greater the electronegativity.
58
What is the relationship between the number of principal energy levels and electronegativity?
The fewer the principal energy levels, the greater the electronegativity.
59
What happens to electronegativity as we move up and to the right on the periodic table?
Electronegativity increases ## Footnote This trend is due to smaller atomic radii and higher nuclear charge.
60
Which atom is the most electronegative?
Fluorine (F) ## Footnote Fluorine has a small atomic radius and a high nuclear charge.
61
What is a polar bond?
A covalent bond in which there is an unequal sharing of the electrons due to differing electronegativities of the atoms involved ## Footnote In polar bonds, electron density is distributed unequally.
62
What does the dipole in a polar bond indicate?
The presence of two poles, one slightly positive and one slightly negative ## Footnote This is represented using partial charges (δ+ and δ-).
63
In the H-F bond, why is the electron density distributed unequally?
Because the nuclei are not identical ## Footnote This leads to a polar bond between hydrogen and fluorine.
64
What is represented by the Greek symbol 'delta' in polar bonds?
Partial charges ## Footnote δ+ indicates a slightly positive side and δ- indicates a slightly negative side.
65
What is the electronegativity value of Fluorine (F)?
3.0 ## Footnote Fluorine is the most electronegative element on the periodic table.
66
Fill in the blank: The bond in hydrogen fluoride (H-F) is ______.
polar
67
True or False: Electronegativity decreases as you move down a group in the periodic table.
True
68
What does a dot and cross diagram represent?
The distribution of electrons in a covalent bond ## Footnote It visually shows how electrons are shared between atoms.
69
What defines the atomic radius of an element?
The distance from the nucleus to the outermost electron shell ## Footnote A smaller atomic radius typically correlates with higher electronegativity.
70
What are the two sides of a dipole in a polar bond referred to as?
Slightly positive and slightly negative ## Footnote These are indicated with partial charges.
71
What is the significance of shielding in electronegativity?
It affects the effective nuclear charge experienced by outer electrons ## Footnote Greater shielding results in lower electronegativity.
72
Why are symmetrical molecules with polar bonds not polar ?
Dipoles cancel out
73
What type of bonding leads to the formation of giant or simple molecules?
Covalent bonding
74
What is an example of a simple molecule?
Water
75
What holds simple molecules together in their solid states?
Intermolecular forces
76
How do intermolecular forces compare to covalent bonds?
Intermolecular forces are much weaker than covalent bonds
77
What are the three types of intermolecular forces?
* Induced dipole-dipole forces * Permanent dipole-dipole forces * Hydrogen bonding
78
Which intermolecular force is the weakest?
Induced dipole-dipole forces
79
Which intermolecular force is the strongest?
Hydrogen bonding
80
What type of forces exist between all molecules?
Induced dipole-dipole forces
81
What additional forces are present if a molecule is polar?
Permanent dipole-dipole forces
82
What type of bonds need to be present for hydrogen bonding to occur?
N-H, O-H, or F-H bonds
83
How does the size of a molecule affect the strength of intermolecular forces?
Bigger molecules have stronger London dispersion forces
84
What are Induced dipole-dipole forces?
Forces that occur when a temporary dipole in one molecule induces a dipole in another molecule.
85
What is a Temporary Dipole?
A situation where one side of a molecule becomes more negative than another due to the movement of electrons.
86
How does a Temporary Dipole affect neighboring atoms?
It induces a dipole in adjacent molecules, leading to an attraction between positive and negative charges.
87
What is the relationship between molecule size and Induced dipole-dipole forces?
Larger molecules exhibit stronger Induced dipole-dipole forces.
88
What factor increases Induced dipole-dipole forces between molecules?
Increased surface area contact.
89
True or False: All simple molecules have Induced dipole-dipole intermolecular forces.
True.
90
How do Induced dipole-dipole forces influence physical properties?
They affect the melting points and boiling points of substances.
91
Fill in the blank: The larger the molecule, the ________ the intermolecular forces.
stronger
92
Which molecule is bigger: iodine or bromine?
Iodine.
93
Which molecule has stronger Induced dipole-dipole intermolecular forces: iodine or bromine?
Iodine.
94
Why does I2 have a higher melting point than Br2?
I2 is a larger molecule with stronger London forces.
95
What are Permanent Dipole-Dipole Forces?
Forces between polar molecules
96
What characterizes a polar bond?
Difference in electronegativity between the atoms
97
What occurs between the δ+ on one molecule and the δ- on an adjacent molecule?
Attraction
98
Provide an example of permanent dipole-dipole forces.
Between HCl molecules
99
In HCl, what does the δ+ on one molecule attract?
The δ- on another molecule
100
What causes a polar bond in HCl?
Difference in electronegativity between H and Cl
101
What is the effect of permanent dipole-dipole forces on the overall strength of intermolecular forces?
Increases the overall strength
102
What remains the overriding influence on the size or molecular weight of a molecule despite the presence of permanent dipole-dipole forces?
The size or Mr of the molecule
103
Fill in the blank: The presence of a permanent dipole-dipole force increases the overall strength of the _______.
IMFs
104
True or False: Permanent dipole-dipole forces occur only between non-polar molecules.
False
105
What is hydrogen bonding?
An attractive force between molecules that have a hydrogen directly bonded to oxygen, nitrogen, or fluorine. ## Footnote Hydrogen bonding is not a true bond but a strong intermolecular force.
106
What types of molecules can participate in hydrogen bonding?
Polar molecules with hydrogen bonded to O, N, or F. ## Footnote O, N, and F are the most electronegative atoms.
107
Why does hydrogen bonding occur specifically with O, N, and F?
These atoms have a big difference in electronegativity with hydrogen. ## Footnote This difference leaves the hydrogen nucleus with a partial positive charge.
108
What effect does electronegativity have on hydrogen in a hydrogen bond?
It pulls the pair of electrons in the bond strongly towards the electronegative atom. ## Footnote This creates a partial positive charge on hydrogen.
109
What is attracted to hydrogen in hydrogen bonding?
A lone pair of electrons on a neighboring molecule containing oxygen, nitrogen, or fluorine. ## Footnote This attraction creates the hydrogen bond.
110
List three features of a hydrogen bonding diagram.
* Partial charges on all atoms * All lone pairs clearly shown * Hydrogen bond clearly shown between the lone pair and the delta + H on the other molecule. ## Footnote These features help visualize the nature of hydrogen bonds.
111
True or False: Hydrogen bonding is a type of covalent bond.
False ## Footnote Hydrogen bonding is an attractive force, not a bond.
112
Fill in the blank: Hydrogen bonding occurs between molecules with hydrogen bonded to ______.
[oxygen, nitrogen, or fluorine]
113
Provide an example of hydrogen bonding between two molecules.
Hydrogen bond between two molecules of NH3. ## Footnote The hydrogen bond is formed between the lone pair on nitrogen and the hydrogen of another NH3 molecule.
114
Provide another example of hydrogen bonding.
Hydrogen bond between two molecules of water (H2O). ## Footnote Each water molecule can form multiple hydrogen bonds due to its two hydrogen atoms and lone pairs on oxygen.
115
What is the reason why ice is less dense than liquid water?
Ice has water molecules that are further apart than in liquid water, resulting in more gaps between the molecules.
116
What is the consequence of ice being less dense than liquid water?
Ice floats on liquid water.
117
What explains the relatively high melting point and boiling point of water?
The presence of hydrogen bonds contributes to the high melting and boiling points.
118
Fill in the blank: The boiling point of water is _______ °C.
100
119
What trend is observed when comparing the hydrides of group 6 elements?
Anomalously high boiling points due to hydrogen bonds.
120
True or False: The strength of intermolecular forces increases with molecular mass.
True
121
What are the intermolecular forces present in water that contribute to its unique properties?
Hydrogen bonds
122
Explain why water has a higher boiling point compared to other group 6 hydrides.
Due to the presence of hydrogen bonds which are stronger than other types of intermolecular forces.
123
Fill in the blank: The density of ice is _______ than that of liquid water.
less

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