Topic 2 - Bonding Flashcards
(135 cards)
1
Q
what is ionic bonding?
A
the strong electrostatic attraction between oppositely
charged ions
2
Q
what are 2 factors that can affect the strength of an ionic bond?
A
ionic radii
ionic charges
3
Q
what affect does ionic charge have on an ionic bond?
A
the greater the charge on an ion, the stronger the ionic bond and therefore there is a higher melting/boiling point
4
Q
what effect does ionic radii have on an ionic bond?
A
Smaller ionic radii allow ions to be closer together, increasing electrostatic attraction and forming stronger ionic bonds.
Larger ionic radii increase the distance between ions, decreasing electrostatic attraction and weakening the bond.
5
Q
what are cations?
A
positive ions
6
Q
what are anions?
A
negative ions
7
Q
Why do atoms gain or lose electrons?
A
to achieve a full outer electron shell and become more stable, forming ions in the process
8
Q
do metal ions form anions or cations?
A
cations - they lose electrons to become positive ions
9
Q
do non-metal ions form anions or cations?
A
anions - they gain electrons to become negative ions
10
Q
What happens to electrons during ionic bonding?
A
Electrons are transferred from the metal atom to the non-metal atom.
11
Q
What is the result of electron transfer in ionic bonding?
A
Both atoms form stable ions with full outer electron shells.
12
Q
Elements in group 1 form________
Elements in group 2 form________
Elements in group 6 form________
Elements in group 7 form________
A
1+ ions
2+ ions
2 - ions
1 - ions
13
Q
what is the trend of ionic radii as you go down a group?
A
ionic radii increases down the group as the ions have more shells of electrons
14
Q
what are isoelectronic ions?
A
ions of different atoms with the same number of electrons
15
Q
what is the trend of ionic radius in isoelectronic ions?
A
decreases as the atomic number increases
16
Q
what is the structure of an ionic compound?
A
giant ionic lattice structure
17
Q
what are the physical properties of an ionic compound?
A
high melting/boiling point
solid at room temperature
can conduct electricity when molten or in solution
soluble in water
18
Q
what is the evidence for the existence of ions?
A
when you electrolyse a green solution of copper (II) chromate (VI) on a pice of wet filter paper, the filter paper will turn blue at the cathode (negative electrode) and yellow at the anode (positive electrode)
Copper (II) ions are blue in solution and chromate (VI) ions are yellow
copper (II) chromate (VI) solution is green as it contains both ions
when you pass a current through the solution the positive ions move to the cathode and the negative ions move to the anode
19
Q
what is a covalent bond?
A
the strong electrostatic force of attraction between 2 nuclei and the shared pair of electrons between them
20
Q
how are molecules formed?
A
when 2 or more atoms bond together and are held together by covalent bonds
21
Q
what is bond energy (bond enthalpy)?
A
the energy required to break one mole of a particular covalent bond in a gaseous state
22
Q
how is bond strength measured?
A
bond energy
23
Q
what is bond length?
A
the distance between the nuclei of 2 atoms that are joined together by a covalent bond - where the attractive and repulsive forces balance
24
Q
what is the relationship between bond strength and bond length?
A
stronger bonds tend to be shorter in length due to the increased electrostatic attraction between the shared pair/s of electrons and the nuclei, pulling the atoms closer together
25
what attracts the positive nuclei in covalent molecules?
the area of electron density (the shared electrons) between the 2 nuclei
26
what must be maintained for a stable covalent bond?
A balance between the attractive and repulsive forces
27
what causes repulsion in a covalent bond
both the negatively charged electron pairs repelling each other and the positively charged nuclei of the bonded atoms repelling each other slightly causes repulsion although this is balanced by the attraction between nuclei and shared electrons that holds the bond together
28
how does electron density affect bond strength?
higher electron density increases the attraction between atoms, strengthening the bond
29
how does bond strength relate to bond length and bond enthalpy?
stronger bonds have higher bond enthalpy and shorter bond lengths
30
why does higher electron density lead to shorter bond lengths?
because more attraction pulls the nuclei closer together
31
what is a dative covalent bond?
it forms when the shared pair of electrons in the covalent bond only come from one bonding atom (can also be called co-ordinate bonding)
32
what determines the shape of a molecule or ion?
number and arrangement of electron pairs surrounding its central atom
33
what are bonding pairs?
bonding pairs are involved in forming covalent bonds with other atoms
34
what are lone pairs?
lone pairs are not involved in bonding and they remain on the central atom
35
what is the electron pair repulsion theory?
a molecule adopts a shape that minimises the repulsion between its electron pairs due to their negative charge
36
what type of electron pair repels the most?
lone pairs
37
what is the order of the angle that electron pairs repel from biggest to smallest?
lone pair-lone pair
lone pair-bonding pair
bonding pair-bonding pair
(Lone pair-lone pair > Lone pair-bonding pair > Bonding pair-bonding pair)
38
when a central atom has 2 electron pairs both of which are bonding pairs, what is the shape and bond angle?
linear shape
180*
39
when a central atom has 3 electron pairs all of which are bonding pairs what is the molecular shape and bond angle?
trigonal planar
120*
40
when the central atom has 4 electron pairs all of which are bonding pairs, what is molecular shape and bond angle?
tetrahedral
109.5*
41
when the central atom has 4 electron pairs 3 of which are bonding pairs and one is a lone pair, what is molecular shape and bond angle?
trigonal pyramidal
107*
42
when the central atom has 4 electron pairs 2 of which are bonding pairs and 2 are lone pairs, what is molecular shape and bond angle?
bent (non-linear)
104.5*
43
when the central atom has 5 electron pairs all of which are bonding pairs, what is molecular shape and bond angle?
trigonal bipyramidal
120° between atoms in the trigonal (equatorial) plane
90° between atoms in the axial and equatorial positions
44
when the central atom has 5 electron pairs 4 of which are bonding pairs and 1 is a lone pair, what is molecular shape and bond angle?
seesaw
<120° in the equatorial positions
<90° between axial and equatorial positions
(The angles are slightly less than in a perfect trigonal bipyramidal shape due to lone pair repulsion.)
45
when the central atom has 5 electron pairs 3 of which are bonding pairs and 2 are lone pairs, what is molecular shape and bond angle?
T-Shaped
<90*
46
when the central atom has 6 electron pairs all of which are bonding pairs, what is molecular shape and bond angle?
octahedral
90* between all bonded atoms
47
when the central atom has 6 electron pairs 4 of which are bonding pairs and 2 are lone pairs, what is molecular shape and bond angle?
square planar
90*
48
what is the shape and bond angle for BeCl2
linear
180*
49
what is the shape and bond angle for BCl3?
trigonal planar
120*
50
what is the shape and bond angle for CH4?
tetrahedral
109.5*
51
what is the shape and bond angle for NH4+?
tetrahedral
109.5*
52
what is the shape and bond angle for H2O?
bent (non-linear)
104.5*
53
what is the shape and bond angle for CO2?
linear
180*
(a double bond is present nut does not affect)
54
what is the shape and bond angle for PCl5 (g)?
trigonal bipyramidal
120* and 90*
55
what is the shape and bond angle for SF6 (g)?
octahedral
90* for all angles
56
predict the shape of OF2
Step 1: Identify the central atom
Oxygen (O) is the central atom.
Step 2: Calculate total outer shell electrons
Oxygen is in group 6, hence it has 6 outer shell electrons
Step 3: Add electrons for each fluorine atom
Each fluorine atom contributes one electron, adding 2 electrons for 2 fluorines
Total electrons = 6 (from O) + 2 (from 2 F) = 8 electrons
Step 4: Calculate total electron pairs
Step 5: Deduce bonding and lone pairs
2 fluorine atoms mean 2 bonding pairs, so 4 - 2 = 2 lone pairs.
Step 6: Predict the molecular shape
Diagram showing the non-linear shape of the OF2 molecule with a bond angle of 104.5 degrees.
With 2 bonding and 2 lone pairs, OF2 has a bent or non-linear shape with an approximate bond angle of 104.5°.
57
Name 3 examples of substances with giant covalent structures
Diamond
Graphite
Silicon dioxide (SiO₂)
58
What are the key properties of giant covalent structures?
Very high melting and boiling points
Insoluble in most solvents
Do not conduct electricity (except graphite)
Very hard (except graphite, which is soft/slippery)
59
Why does diamond not conduct electricity?
All four outer electrons of each carbon atom are used in strong covalent bonds, so there are no free electrons to carry charge.
60
What makes graphite soft and slippery?
Graphite’s layers are held together by weak van der Waals forces, allowing them to slide over each other easily.
61
what is electronegativity?
the ability of an atom to attract the bonding electrons in a covalent bond
62
what are properties of more electronegative elements?
higher nucleur charges
smaller atomic radii
(therefore electronegativity increases across periods and up the groups - not noble gases)
63
what is shielding?
Shielding is the reduction in attraction between the nucleus and the outer electrons caused by inner electron shells repelling the outer electrons
64
what does is mean if a compound containing elements of similar electronegativity - a small electronegativity difference?
it is purely covalent
65
what does is mean if a compound containing elements of very different electronegativity - a large electronegativity difference?
it is ionic
66
how is a permanent dipole (polar covalent bond) formed?
when, the elements in the bond have different electronegativities, the bonding electrons are more strongly attracted to the more electronegative atom, this unequal sharing of electrons makes the bond polar - polar bonds have a permanent electric dipole
67
what is a dipole?
A dipole is created when a molecule or bond has a separation of charge due to a difference in electronegativity between two atoms.
One end becomes slightly negative (δ–) and the other slightly positive (δ+)
68
what are the 4 types of intermolecular forces in order of increasing strength?
London forces
dipole-induced dipole forces
dipole-dipole forces
hydrogen bonding
69
what do van Der Waals' forces include?
London forces
dipole-induced dipole forces
dipole-dipole forces
70
where do London forces occur?
between all molecules and noble gases
71
where do dipole-induced dipole forces occur?
between polar and non-polar molecules
72
where do dipole-dipole forces occur?
between all polar molecules
73
where does hydrogen bonding occur?
between polar molecules with H-F, H-O or H-N bonds
74
what is another name for London forces?
instantaneous, induced dipole-dipole interactions
75
how do London forces occur?
electrons in atoms are constantly moving, so at any moment there may be more electrons on one side of the atom than the other creating a temporary (instantaneous) dipole
this temporary dipole can induce an opposite dipole in a neighbouring atom causing a weak electrostatic attraction. between the atoms
this induced dipole can then induce further dipoles in other nearby particles
although these dipoles are constantly forming and disappearing as the electrons move the overall effect is a net attraction between that atoms or molecules
76
what are the factors affecting the strength of London forces?
size - larger atoms and molecules gave more electrons and a greater volume of electron density that can become polarised creating stronger temporary dipoles
polarisability - atoms or molecules with more easily distorted electron clouds (higher polarisability) experience stronger induced dipole-dipole forces. polarisability generally increases with the size and number of electrons in an atom or molecule
77
substances with stronger London forces tend to have _______ boiling points
higher
78
what are London forces strong enough to hold?
molecules together in a lattice structure
79
how do London forces explain the existence of noble gas liquids and solids?
even though noble gas atoms have completer outer shells and do not form covalent, ionic or metallic bonds, the weak London forces allow them to condense into the liquid and solid states at very low temperature
80
what is a metallic bond?
the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons
81
what force do polar molecules experience?
permanent dipole-dipole forces
82
how do permanent dipoles arise?
from the unequal sharing of electrons in covalent bonds, the partial positive and partial negative charges on polar molecules enable them to experience permanent dipole-dipole forces
83
what are permanent dipole-dipole forces?
the electrostatic attractions between the partial positive end of one polar molecule and the partial negative end of another
84
What type of intermolecular forces exist between all molecules?
Induced dipole–dipole forces (London forces) exist between all molecules, whether polar or non-polar
85
What extra type of intermolecular force do polar molecules have?
Polar molecules also have permanent dipole–dipole forces, in addition to London forces.
86
Why do polar molecules have stronger intermolecular forces than non-polar molecules of a similar size?
Because polar molecules experience both London forces and permanent dipole–dipole forces, giving them stronger overall intermolecular forces
87
What effect do stronger intermolecular forces have on boiling points?
Stronger intermolecular forces mean more energy is needed to separate molecules, so polar molecules usually have higher boiling points than non-polar ones of similar size.
88
what is hydrogen bonding?
a type of (permanent) dipole-dipole force that occurs when hydrogen is bonded to the highly electronegative elements fluorine, oxygen or nitrogen
89
what is needed for hydrogen bonding to occur?
the molecule must contain a hydrogen atom covalently bonded to either F, O or N as they are highly electronegative
there must be a lone pair of electrons on the F, O or N atom of an adjacent molecule available to interact with the hydrogen
90
how do hydrogen bonds form?
the H-F, H-O and H-N bonds are highly polar due to the large electronegativity differences between hydrogen and these elements, this leads to a significant partial positive charge on the H atom and a partial negative charge on the F,O or N atom
the positively charged hydrogen is strongly attracted to the negatively charged lone pair on the F, O or N atom forming a hydrogen bond between molecules
91
give 3 examples of molecules contain H-N, then O-H and then H-F
ammonia (NH3)
water (H2O)
hydrogen fluoride (HF)
92
what are properties of substances containing hydrogen bonds?
greater solubility in water - substances that can form hydrogen bonds with water tend to be soluble while those that cannot are typically insoluble
higher melting and boiling pints compared to similar-sized molecules that cannot hydrogen bond as extra energy is needed to overcome the strong hydrogen bonding forces
93
how does hydrogen bonding explain the fact that ice is less dense than water?
In solid ice, water molecules are arranged in a regular 3D lattice, This structure is held together by hydrogen bonds, which keep the molecules further apart than they would be in liquid water Hydrogen bonds are longer than covalent bonds, so they create open spaces in the structure.
This makes ice take up more space (volume) but with the same mass, so it's less dense
Some hydrogen bonds break, and water molecules can move closer together.
This makes liquid water denser than ice, which is why ice floats on water.
94
in terms of intermolecular forces explain the trends in boiling temperatures of alkanes with increasing chain length
as chain length increases, boiling point increases, as there are stronger London forces, and there is greater surface area in straight chain alkanes as there is more contact between molecules and therefore stronger intermolecular attractions, so there is more energy needed to overcome the stronger London forces so there is a higher boiling point
95
in terms of intermolecular forces explain the effect of branching in the carbon chain on the boiling temperatures of alkanes
As branching in the carbon chain increases, the boiling point of alkanes decreases
branched alkanes are more compact and spherical in shape, giving them a smaller surface area in contacts with neighbouring molecules, so there are fewer London forces between molecules and less energy is needed to overcome the weaker intermolecular forces and therefore there is a lower boiling point
96
in terms of intermolecular forces explain the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
alcohols contain an -OH group (hydroxyl) which allows them to form hydrogen bonds between molecules, hydrogen bonding is much stronger than the London forces found in alkanes, the hydrogen bonds require more energy to break and therefore alcohols have higher boiling points and lower volatility than alkanes in terms of intermolecular forces explain the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
97
in terms of intermolecular forces explain the trends in boiling temperature of the hydrogen halides HCl to HI
The boiling points of hydrogen halides generally increase from HCl to HI as the number of electrons increase so the London forces become stronger so the boiling points increase as you go down the group
98
why does HF have an unusually high boiling point compared to the rest of the hydrogen halides?
HF has hydrogen bonding which is a lot stronger than the London forces found in the other hydrogen halides and therefore it has a higher boiling point
99
what 3 key processes must occur in order for a substance to dissolve in a solvent?
Bonds within the substance (solute) molecules must break.
Bonds between the solvent molecules must break.
New bonds must form between the solute and solvent molecules
100
what is a solvent?
a substance that dissolves a solute to form a solution
101
on what condition does a substance usually dissolve?
usually a substance will only dissolve if the strength of the new bonds formed is about the same as or greater than the strength of the bonds that are broken
102
what are the 2 main types of solvents?
polar solvents
non-polar solvents
103
what else can water be referred to as?
an aqueous solvent
104
ionic substances often dissolve in what kind of solvent?
polar solvents like water
105
in which process do ionic substances often dissolve in solvents?
hydration
106
describe the process of hydration
water molecules are polar with δ+ hydrogen atoms and δ- oxygen atoms
when an ionic substance is added to water the ions are attracted to the oppositely charged ends of the water molecules
the water molecules surround and separate the ions from the ionic lattice
107
why do some ionic compounds like aluminium oxide (Al2O3) not dissolve in water
because the ionic bonds are stronger than the potential bonds with water molecules
(for Al2O3 - its because the high charge density of the Al3+ ion leads to strong ionic bonding)
108
why are alcohols generally soluble in water?
due to their ability to form hydrogen bonds
109
how are alcohol's dissolved in polar solvents such as water?
the polar O-H bond in an alcohol is attracted to the polar O-H bonds in water. Hydrogen bonds form between the lone pairs on the δ- oxygen atoms and the δ+ hydrogen atoms
110
what affect does the length of the carbon chain have on the solubility of the alcohol
the more carbon atoms there are the less soluble the alcohol is as the carbon chain is not attracted to water
111
why do halogenoalkanes not readily dissolve in water?
the dipoles in halogenoalkanes are not wrong enough to form hydrogen bonds with water molecules
the hydrogen bonding between water molecules is stronger than the potential bonds with halogenalkanes preventing dissolution
112
what solvents will halogenoalkanes dissolve in?
halogenoalkanes can form permanent dipole-dipole bonds and thus dissolve in polar solvents that also form such interactions rather than hydrogen bonds
113
when do non-polar substances dissolve best?
in non-polar solvents
114
why do non-polar substances dissolve best in non polar solvents?
non-polar molecules have London forces between them, they form similar London forces with non-polar solvents allowing dissolution
115
why do non-polar substances generally not dissolve well in water?
water molecules have stronger attractions to each other than to non-polar molecules so they don't dissolve well
116
when do substances dissolve best?
in solvents with similar intermolecular forces
117
what are allotropes?
the different structural forms of an element in the same physical state
118
what are the 3 allotropes of carbon?
diamond
graphite
graphene
119
what is the diamond structure?
Each carbon atom forms 4 very strong covalent bonds with others in a tetrahedral arrangement
120
what are the properties of diamond?
Extremely hard - Extensive network of strong covalent bonds not easily broken.
Very high melting point - Huge amount of energy needed to break enough bonds to melt diamond.
Good thermal conductor - Strong interatomic bonds transmit heat through vibrations.
Electrical insulator - All outer electrons tied up in localised bonds so no free electrons to carry charge.
Insoluble - Covalent bonds too strong to be broken by solvation
121
what is the structure of graphite?
Each carbon atom forms 3 strong covalent bonds in a planar hexagonal pattern, with each carbon contributing 1 delocalised electron.
Multiple stacked layers of hexagonal carbon arrays with weak intermolecular forces between layers.
122
what are the properties of graphite?
Soft and slippery - Weak intermolecular forces let sheets slide over each other.
Conducts electricity along layers - Delocalised electrons move through the 2D lattice carrying electrical charge.
Lower density than diamond - Weak intermolecular forces lead to increased separation between layers.
High sublimation temperature but lower melting point than diamond - Covalent bonds within each layer are very strong but the weaker intermolecular forces between layers means graphite melts at a lower temperature.
123
what is the structure of graphene?
Graphene consists of a single layer of carbon atoms interconnected through strong planar covalent bonds in a hexagonal pattern, with each carbon contributing 1 delocalised electron. This essentially forms a one-atom thick slice of graphite.
124
what are the properties of graphene
Excellent electrical and thermal conductivity - Delocalised electrons move through the 2D lattice transporting heat and charge.
Very strong - Extensive network of covalent bonds not easily broken.
Transparent and extremely lightweight - A single layer of atoms light and thin enough to transmit visible light.
125
describe the structure of a giant metallic lattice
electrons in the outermost shell of the metal atoms are delocalised leaving a positive metal ions, the positive metal ions are electrostatic ally attracted to the delocalised negative electrons, they form a lattice of closely packed positive ions in a sea of delocalised electrons
126
when does metallic bonding occur?
between metal cations and delocalised electrons
127
what is the strength of a metallic bond influenced by?
number of delocalise electrons per atom
charge of the metal cation
radius of the metal cation
128
why does the Number of delocalised electrons per atom affect the strength of metallic bonds?
Metals with a higher number of delocalised electrons per atom tend to form stronger metallic bonds.
The increased number of delocalised electrons allows for stronger electrostatic attractions between the electrons and the metal cations.
129
why does the charge of the metal cation affect the strength of metallic bonds?
A higher cation charge results in stronger electrostatic attractions between the cation and the delocalised electrons
130
why does the radius of the metal cation affect the strength of metallic bonds?
Smaller metal cations have a higher charge density, which allows them to hold the delocalised electrons closer to the nucleus.
This proximity enhances the electrostatic attractions between the cation and the electrons, resulting in a stronger metallic bond
131
what are examples of covalently bonded substances with simple molecular structure?
iodine - I2
Ice
H2O - water
132
why do metals have high melting and boiling points?
The strong electrostatic forces of attraction between the positively charged metal cations and the sea of delocalised electrons must be overcome for the lattice to break apart; this requires large amounts of energy
133
why are metals good conductors of electricity and heat?
The delocalised electrons can flow freely through the lattice to transfer charge and heat energy
134
what are metals malleable and ductile?
Layers of the cation lattice can slide over one another when the metal is hammered or pulled because there are no bonds locking individual cations together
135
why are metals insoluble?
The strength of the metallic bonding prevents water or other solvent molecules from pulling the cations away from the lattice structure and dissolving the metal
