2.2.2 Bonding and structure Flashcards
1
Q
what is an ionic bond definition?
A
ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer
2
Q
when is ionic bonding stronger?
A
when the ions are smaller and/or have higher charges
3
Q
what pattern are the ions in an ionic solid arranged in?
A
a giant ionic lattice
4
Q
why do ionic compounds have high melting points?
A
there are strong electrostatic attractive forces between the oppositely charged ions in the lattice
which means a lot of energy is required to break these bonds
5
Q
why do ionic compounds not conduct electricity when in a solid?
A
the ions are held together tightly in the lattice and cannot move so no charge is conducted
6
Q
why are ionic compounds good conductors of electricity when in a solution or molten?
A
the ions are free to move when molten or in a solution so a charge can be carried
7
Q
why are ionic compounds soluble in water?
A
because water molecules have a slight electrical charge so can attract the ions away from the lattice
8
Q
what can melting points in ionic compounds be affected by?
A
nuclear charge
ionic radius
9
Q
what is the definition for covalent bonding?
A
a covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
10
Q
why do simple covalent compounds have low melting and boiling points?
A
because there are weak intermolecular forces between the covalent bonds so little energy is needed to break the bonds
11
Q
what state are simple covalent compounds usually in at room temperature?
A
usually gases or volatile liquids due to the weak intermolecular forces
12
Q
why do simple covalent compounds not conduct electricity?
A
there are no free ions or electrons present in a simple covalent compound so they do not conduct electricity
13
Q
why do giant covalent compounds have high melting and boiling points?
A
because they have strong covalent bonds in the giant molecule so a lot of energy is needed to break these apart
14
Q
examples of giant covalent compounds
A
diamond
graphite
silicon dioxide
15
Q
why does graphite conduct electricity?
A
because it has delocalised electrons that can move throughout the whole structure and carry charge
16
Q
when is a dative covalent bond formed?
A
when one atom contributes both of the electrons needed for the covalent bond
17
Q
what is needed for a dative covalent bond to form?
A
one atom has to have a lone pair of electrons and the other atom must have a vacant orbital
18
Q
how is a dative covalent bond represented in a drawing?
A
an arrow which shows the direction of the electron pair donation
19
Q
what is the definition of a metallic bond?
A
the electrostatic force of attraction between a lattice of positively charged ions and free/delocalised electrons
20
Q
what are the factors that affect the strength of a metallic bond?
A
size of the charge on the positive ions - the bigger this is the stronger the bond
the number of mobile electrons - the more there are, the stronger the bond
21
Q
why are metallic compounds good conductors of electricity?
A
because they have delocalised electrons
22
Q
why do metallic compounds have high melting and boiling points?
A
because they have a giant lattice structure and a strong electrostatic attraction between the delocalised electrons and the metal ions which means a lot of energy is required in order to break the bonds
23
Q
what is the definition of electronegativity?
A
the ability of an atom to attract the pair of electrons in a covalent bond (electron density)
24
Q
what scale is electronegativity measured on?
A
the Pauline scale
25
what element is the most electronegative?
fluorine (4.0)
26
what element is the least electronegative?
francium (0.7)
27
which direction on the periodic table does the electronegativity gets weaker and weaker?
down and across from fluorine
28
what 2 factors affect the electronegativity of an atom?
the size of the nuclear charge (number of protons)
the size of an atom
29
how does the size of nuclear charge affect the electronegativity of an atom?
the bigger the nuclear charge is the larger the attraction between the nucleus and the pair of electrons - the electronegativity goes up
30
how does the size of the atom affect the electronegativity of an atom?
as the size increases, the pair of electrons are further from the nucleus, and there will be shielding effect from the inner electron, the electronegativity goes down
31
if both atoms have equal electronegativity where will the electrons be?
the bond is formed roughly half way between the 2 atoms
this is the case in non polar covalent bonds
32
where is the electron if the atoms have slightly different electronegativities?
the electron pair will be pulled towards the more electronegative atom
this end will have more electron density and will become slightly negatively charged
33
what does the polarity of a molecule depend on?
its shape
34
is a molecule with overall symmetry polar or non polar?
non polar
35
is a non symmetrical molecule polar or non polar?
polar
36
what are the weakest type of intermolecular forces?
London forces
37
what do London forces form between?
between neighboring non polar forces
38
what is the distance that London forces have an affect over?
a couple of nanometers
39
how is a temporary dipole formed?
an electron is always on the move
a temporary dipole is formed when at any one time one end of the molecule might have a lower electron density than the other
40
what symbol do you need when drawing an induced dipole?
party drawn 8 with a minus sign in top right hand corner
41
how is an induced dipole formed?
the 8- end of the molecule will attract electron density to one end of the neighboring molecule
the neighboring molecule now has a dipole that has been induced by the first molecule - an induced dipole
42
if the molecule is larger, is the London force stronger or weaker?
stronger
43
why is the London force stronger if the molecule is bigger?
more surface contact
more electrons lead to stronger London forces
44
how much stronger are permanent dipole - dipole forces than London forces?
10 times stronger than London forces
45
where do permanent dipole - dipole forces form between?
between two polar covalent molecules (i.e. in molecules with a difference in electronegativity)
46
do compounds with only London forces have a higher or lower boiling point than compounds that also have permanent dipole - dipole forces?
lower
47
where does hydrogen bonding only happen?
only happens when a hydrogen atom is covalently bonded to a N, O or F atom
48
which is the strongest intermolecular force of attraction?
hydrogen bonding
49
how much stronger is hydrogen bonding than London forces?
100 times stronger
50
what do you need to include when drawing hydrogen bonding?
always show lone pairs
must always show the polarity of the element (8-)
hydrogen must always be from lone pair
51
what does the shape of a molecule depend on?
the number of electron pairs around the central atom in the molecule
52
what is the molecule shape called with 2 bond pairs and no lone pairs?
linear
53
what is the bond angle in a linear molecule?
180 degrees
54
what is the molecule shape called with 3 bond pairs and no lone pairs?
trigonal planar
55
what is the bond angle in a trigonal planar shaped molecule?
120 degrees
56
what is the molecule shape called with 4 bond pairs and no lone pairs?
tetrahedral
57
what is the bond angle in a tetrahedral shaped molecule?
109.5 degrees
58
what is the molecule shape called with 5 bond pairs and no lone pairs?
trigonal bipyramidal
59
what are the bond angles in a trigonal bipyramidal shaped molecule?
120 degrees and 90 degrees
60
what is the molecule shape called with 6 bond pairs and no lone pairs?
octahedral
61
what is the bond angle in an octahedral shaped molecule?
90 degrees
62
do lone pairs or bonded pairs repel more strongly?
lone pairs
63
what is the molecule shape called with 3 bond pairs and 1 lone pair?
pyramidal
64
what is the bond angle in a pyramidal shaped molecule?
107 degrees
65
what is the molecule shape called with 2 bond pairs and 2 lone pairs?
V - shaped
66
what is the bond angle in V-shaped molecule?
104.5 degrees
67
what is the method for working out the shape of an ion?
1 - draw the outer shell electrons of the central atom
2 - workout the charge and remove or gain electrons depending on the charge
3 - pair up the electrons in the usual way
4 - workout the shape and bond angle(s) from the number of bonded pairs and lone pairs
68
why does water have a high melting point and boiling point for being a simple molecule?
because of its hydrogen bonding
69
why is ice less dense than water?
ice has an open lattice with hydrogen bonds holding the water molecules apart, when ice melts the rigid hydrogen bonds collapse, allowing the water molecules to move close together
the distance the molecules are held apart is what makes ice less dense than water
70
why does iodine have a low melting point?
each iodine atom bonds covalently to another to form an I2 molecule
thee molecules then bond to each other through weak intermolecular bonds (London forces)
a low amount of energy is needed to break these bonds
71
What is ionic bonding?
The electrostatic force of attraction between oppositely charged ions formed by electron transfer
## Footnote
Metal atoms lose electrons to form positive ions and non-metal atoms gain electrons to form negative ions.
72
What happens to magnesium (Mg) during ionic bonding?
Mg loses two electrons to form Mg2+
## Footnote
Mg goes from 1s2 2s2 2p6 3s2 to Mg2+ 1s2 2s2 2p6.
73
What happens to oxygen (O) during ionic bonding?
O gains two electrons to form O2-
## Footnote
O goes from 1s2 2s2 2p4 to O2- 1s2 2s2 2p6.
74
How does the size and charge of ions affect ionic bonding strength?
Stronger ionic bonding and higher melting points occur with smaller ions and/or higher charges
## Footnote
E.g., MgO has a higher melting point than NaCl due to smaller and higher charged ions.
75
What is a giant ionic lattice?
A regular 3D pattern in which ions in an ionic solid are arranged
## Footnote
The arrangement is represented diagrammatically but does not show the ionic bonds.
76
What are the typical physical properties of ionic compounds?
- High melting points
- Non conductor of electricity when solid
- Good conductor of electricity when in solution or molten
- Usually soluble in aqueous solvents
## Footnote
High melting points are due to strong electrostatic forces between ions.
77
What is a covalent bond?
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
## Footnote
Covalent bonds involve the sharing of electron pairs between atoms.
78
What is a dative covalent bond?
A bond formed when the shared pair of electrons comes from only one of the bonding atoms
## Footnote
Also called coordinate bonding.
79
What is the shape of the NH4+ ion?
Tetrahedral
## Footnote
The dative covalent bond behaves like an ordinary covalent bond in determining shape.
80
What does the direction of the arrow in a dative covalent bond indicate?
The direction goes from the atom providing the lone pair to the atom that is deficient
## Footnote
This helps visualize the source of the shared electron pair.
81
What is average bond enthalpy?
A measurement of covalent bond strength
## Footnote
The larger the average bond enthalpy value, the stronger the covalent bond.
82
What is ionic bonding?
Electrostatic force of attraction between oppositely charged ions
## Footnote
Examples include sodium chloride and magnesium oxide.
83
What is covalent bonding?
Shared pair of electrons
## Footnote
This bonding forms simple molecular substances.
84
What are examples of simple molecular substances?
* Iodine
* Ice
* Carbon dioxide
* Water
* Methane
## Footnote
These substances have intermolecular forces such as induced dipole-dipole, permanent dipole-dipole, and hydrogen bonds.
85
What is the general property of giant ionic structures regarding boiling and melting points?
High due to strong electrostatic forces between oppositely charged ions
## Footnote
The giant lattice of ions contributes to these high temperatures.
86
How do simple molecular substances generally behave in terms of solubility in water?
Generally poor due to weak intermolecular forces
## Footnote
Specific types of intermolecular forces include induced dipole-dipole and hydrogen bonds.
87
What is the conductivity of giant ionic structures when solid?
Poor, as ions can't move and are fixed in lattice
## Footnote
This limits the ability to conduct electricity.
88
What is the conductivity of ionic substances when molten?
Good, as ions can move freely
## Footnote
This allows for electrical conductivity.
89
What is the general description of the structure of giant ionic compounds?
Crystalline solids
## Footnote
This refers to their organized lattice structure.
90
What is the general description of molecular (simple) compounds?
Mostly gases and liquids
## Footnote
This indicates their physical states at room temperature.
91
What is the bond angle in a linear molecule?
180 degrees
## Footnote
Linear molecules have no bonding or lone pairs.
92
What is the bond angle in a trigonal planar molecule?
120 degrees
## Footnote
This occurs when there are no lone pairs.
93
What is the bond angle in a tetrahedral molecule?
109.5 degrees
## Footnote
This shape occurs with four bonding pairs and no lone pairs.
94
What is the bond angle in a trigonal pyramidal molecule?
107 degrees
## Footnote
Lone pairs reduce the bond angle from the tetrahedral angle.
95
What is the bond angle in a bent molecule?
104.5 degrees
## Footnote
This shape is influenced by lone pairs of electrons.
96
What is the bond angle in an octahedral molecule?
90 degrees
## Footnote
This occurs when there are six bonding pairs.
97
What is the first step in explaining the shape of a molecule?
State the number of bonding pairs and lone pairs of electrons
## Footnote
This sets the foundation for understanding the molecular geometry.
98
What happens to electron pairs in terms of repulsion?
They repel and try to get as far apart as possible
## Footnote
This principle is crucial for determining molecular shape.
99
What is the effect of lone pairs on bond angles?
Lone pairs repel more than bonding pairs, reducing bond angles by about 2.5 degrees per lone pair
## Footnote
This adjustment is important in determining molecular geometry.
100
What is electronegativity?
The relative tendency of an atom in a covalent bond in a molecule to attract electrons to itself.
## Footnote
F, O, N, and Cl are the most electronegative atoms.
101
How does electronegativity change across a period?
Electronegativity increases across a period due to an increase in the number of protons and a decrease in atomic radius.
102
How does electronegativity change down a group?
Electronegativity decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases.
103
What type of bond forms when elements have similar electronegativities?
A purely covalent bond forms with a small electronegativity difference.
104
What is a polar covalent bond?
A bond formed when elements have different electronegativities, resulting in an unequal distribution of electrons and a charge separation (dipole).
105
In a polar covalent bond, which element has the negative end?
The element with the larger electronegativity.
106
What characterizes a symmetric molecule in terms of polarity?
A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecule are polar.
107
What is the result of individual dipoles in a symmetric molecule?
The individual dipoles cancel out due to the symmetrical shape, resulting in no net dipole moment: the molecule is non-polar.
108
What type of bond is formed with elements of very different electronegativities?
An ionic bond.
109
What are induced dipole-dipole interactions?
Interactions that occur between all molecular substances and noble gases, also called London forces.
110
What causes induced dipoles to form?
Fluctuations in electron density in a molecule can create temporary dipoles that induce dipoles in neighboring molecules.
111
What is the main factor affecting the size of induced dipole-dipole interactions?
The number of electrons in a molecule; more electrons increase the chance of temporary dipoles forming.
112
How does the boiling point of halogens change down the group?
The boiling points increase due to the increasing number of electrons in larger molecules, enhancing induced dipole-dipole interactions.
113
How does molecule shape affect induced dipole-dipole interactions?
Long chain alkanes have a larger surface area for interactions compared to spherical branched alkanes.
114
What are permanent dipole-dipole forces?
Forces that occur between polar molecules, stronger than induced dipole-dipole interactions.
115
What types of bonds commonly exhibit permanent dipole-dipole forces?
Bonds such as C-Cl, C-F, C-Br, H-Cl, and C=O.
116
What is the relationship between permanent dipole-dipole forces and induced dipole-dipole interactions?
Permanent dipole-dipole forces occur in addition to induced dipole-dipole interactions.
117
What are van der Waals' forces?
A term that encompasses both permanent dipole-dipole and induced dipole-dipole interactions.
118
What is hydrogen bonding?
It occurs in compounds with a hydrogen atom attached to nitrogen, oxygen, or fluorine that have an available lone pair of electrons.
## Footnote
Examples include -O-H, -N-H, and -F-H bonds.
119
Which three atoms are most commonly involved in hydrogen bonding?
Nitrogen, oxygen, fluorine
## Footnote
These atoms must have a lone pair of electrons available for hydrogen bonding.
120
How does the electronegativity difference influence hydrogen bonding?
There is a large electronegativity difference between hydrogen and the electronegative atoms.
## Footnote
This difference is crucial for the formation of strong hydrogen bonds.
121
What type of interactions does hydrogen bonding occur alongside?
Induced dipole-dipole interactions
## Footnote
Hydrogen bonding is a stronger type of interaction compared to induced dipole-dipole interactions.
122
What is the effect of hydrogen bonding on the boiling points of certain compounds?
It causes anomalously high boiling points for H2O, NH3, and HF.
## Footnote
The presence of hydrogen bonds significantly increases the boiling points of these substances.
123
What trend is observed in the boiling points from H2S to H2Te?
An increase in boiling point is observed due to increasing induced dipole-dipole interactions.
## Footnote
This trend is attributed to the increasing number of electrons in the molecules.
124
Name some types of compounds that can form hydrogen bonds.
* Alcohols
* Carboxylic acids
* Proteins
* Amides
## Footnote
These compounds contain functional groups capable of hydrogen bonding.
125
How many hydrogen bonds can a water molecule form?
Two hydrogen bonds
## Footnote
Water's oxygen atom is very electronegative and has two lone pairs of electrons.
126
Why does ice have a lower density than liquid water?
The molecules in ice are held further apart due to hydrogen bonding.
## Footnote
This arrangement leads to a lower density compared to liquid water.
127
What type of bonds are present between iodine atoms in a molecular iodine (I2) molecule?
Covalent bonds
## Footnote
These covalent bonds hold the iodine atoms together within the molecule.
128
What type of intermolecular forces hold the crystals of iodine together?
Weak induced dipole-dipole interactions
## Footnote
These interactions are responsible for the structure of iodine crystals.
129
True or False: Hydrogen bonding is stronger than induced dipole-dipole interactions.
True
## Footnote
Hydrogen bonds are a specific type of strong intermolecular force.
130
octahedral drawing
131
tetrahedral drawing
132
trigonal planar drawing
133
trigonal bipyramidal drawing
134
linear drawing
135
hydrogen bonding drawing
