1
Q
What is a covalent bond
A
A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded atoms.
2
Q
Why does sharing electrons form a stable bond?
A
Sharing electrons allows atoms to achieve a lower-energy state, where attractive forces (between nuclei and shared electrons) are balanced by repulsive forces (between nuclei), holding atoms at a fixed distance.
3
Q
What is meant by the octet rule?
A
The octet rule is the tendency of atoms to gain a valence shell with eight electrons, giving a stable noble-gas electron configuration.
4
Q
Why does a covalent bond form at a specific distance between two atoms?
A
Because the system becomes most stable when attractive forces (nuclei ↔ shared electrons) are balanced by repulsive forces (nucleus–nucleus and electron–electron). This balance fixes the atoms at a particular distance (the bond length).
5
Q
On a potential energy vs distance graph for a covalent bond, what does the minimum point represent?
A
The equilibrium bond length, where the system’s potential energy is lowest (most stable) and the net force is zero (attraction = repulsion).
6
Q
Why does the energy rise sharply if atoms get too close?
A
Because repulsive forces dominate, especially nucleus–nucleus repulsion (and electron–electron repulsion), causing the potential energy to increase steeply.
6
Q
Why does the energy decrease as two atoms approach each other initially?
A
As atoms approach, nucleus–electron attraction increases and electron density builds between the nuclei, lowering potential energy; this continues until the energy minimum (equilibrium bond length) is reached.
7
Q
What does the “depth” of the potential energy well represent?
A
The bond energy / bond dissociation enthalpy (energy required to break the bond), equal in magnitude to the energy released when the bond forms.
8
Q
In the H₂ graph shown, what do the labeled regions A, B, C indicate?
A
They show how electron density and forces change as distance changes:
- A (far apart): atoms essentially separate, little interaction
- B (approaching): attraction increases, energy decreases
- C (minimum): stable bond at lowest energy (equilibrium distance)
9
Q
What is meant by “covalent bond forms at the point of lowest energy”?
A
The bonded arrangement is favored because systems naturally move toward minimum potential energy; at the minimum, the bond is stable and requires energy input to break.
10
Q
How do you identify bond length from a potential energy curve?
A
The bond length is the distance at the energy minimum (equilibrium separation)
11
Q
If the atoms are at a distance greater than the equilibrium bond length, what is the net force?
A
Net force is attractive (pulls atoms closer) because attraction dominates over repulsion at larger separations until equilibrium is reached.
12
Q
If the atoms are at a distance shorter than the equilibrium bond length, what is the net force?
A
Net force is repulsive (pushes atoms apart) because nucleus–nucleus and electron–electron repulsions dominate.
13
Q
explanation linking the potential energy curve to stability of a covalent bond.
A
As atoms approach, increasing nucleus–electron attraction lowers potential energy until a minimum is reached where attraction and repulsion balance. This minimum corresponds to the bond length and maximum stability; moving closer increases repulsion sharply, and moving further weakens attraction, so energy rises in either direction.
14
Q
Why is the octet rule a guideline rather than a law? mention exceptions
A
The octet rule is a guideline because the most stable structure is the one that minimises energy and formal charge, which can lead to exceptions:
- Hydrogen (2 electrons only)
- Incomplete octets (Be, B, Al)
- Expanded octets (period 3 and beyond, HL)
- Radicals
- Transition elements (variable electron arrangements)
15
Q
Why do noble gases form covalent bonds less readily than other elements?
A
Noble gases already have a complete valence shell, so they have no energetic advantage in sharing electrons.
16
Q
Why can covalent bonds form between atoms of the same element?
A
Identical atoms have equal electronegativity, so both nuclei attract the shared pair equally, forming a stable bond through nucleus–electron attraction without the large electronegativity difference needed for ionic bonding.
17
Q
What is a Lewis formula (Lewis structure)?
A
A representation showing all valence electrons (bonding and non-bonding) in a covalently bonded species using dots, crosses, or lines.
18
Q
What do lone pairs represent in a Lewis structure?
A
Lone pairs are non-bonding electron pairs localised on one atom; they increase electron-domain repulsion, distort bond angles, and often act as electron-pair donors (Lewis bases), affecting shape and reactivity.
19
Q
Step 2 of drawing a Lewis structure?
A
Draw the skeletal structure, connecting the central atom to surrounding atoms using single bonds.
20
Q
When counting electrons around an atom, do you track where electrons came from?
A
No. You only count how many electrons surround the atom, including all bonding electrons.
20
Q
Step 1 of drawing a Lewis structure?
A
Calculate the total number of valence electrons, adding 1 for each negative charge and subtracting 1 for each positive charge.
21
Q
How do you choose the central atom?
A
The central atom is usually the least electronegative (so it can share electrons with multiple atoms) and the one able to form the most bonds; hydrogen is always terminal because it forms only one bond.
22
Q
Step 3 of drawing a Lewis structure?
A
Add electrons to outer atoms first to complete their octets, then to the central atom.
23
What if there are not enough electrons to complete octets?
Convert lone pairs on outer atoms into double or triple bonds.
- keeps total electron count unchanged and reduces unfavourable formal charges
24
Step 4 of drawing a Lewis structure?
Check that the total number of electrons used matches the calculated total.
25
How many electrons does a single bond contribute to each atom?
Two electrons — bonding electrons are shared and counted fully by both atoms.
26
Why does an oxygen with a single bond still have an octet?
Because it has 6 electrons in lone pairs + 2 electrons from the shared bond = 8.
27
Why does it seem like an electron “appears from nowhere” in negative ions?
Because a negative charge means an extra electron is added to the entire ion, not to a specific atom.
27
How is a negative charge represented in a Lewis structure?
By adding electrons to the total count and enclosing the structure in square brackets with the charge outside.
28
Is the extra electron in NO₃⁻ on one oxygen?
The negative charge is delocalised by resonance: the π electrons and negative charge are spread over all three oxygens, giving three equivalent N–O bonds with equal intermediate bond order.
29
Does the extra electron “move around” in resonance structures?
The electron density is delocalised, meaning no single atom permanently holds the charge.
30
What does a positive charge mean in a Lewis structure?
The species has lost electrons, so electrons are subtracted from the total count.
31
Is the lost electron removed from a specific atom?
No. The electron is removed from the entire molecule, not a specific atom.
31
Why does NO⁺ form a triple bond?
NO⁺ has fewer total valence electrons than NO, so a triple bond is needed to place 8 electrons around both N and O (octets) while using the reduced electron count, giving a stable low-energy structure
32
What is an incomplete octet?
A stable structure where the central atom has fewer than eight electrons in its valence shell, e.g. BeCl₂, BF₃.
33
Why do small atoms form incomplete octets?
Be and B are electron-deficient and stabilise with fewer than 8 electrons because forming additional bonds would require unfavourable charge separation / high formal charges; the lowest-energy Lewis structure often has an incomplete octet.
Be commonly forms two covalent bonds (4 electrons around Be), and B commonly forms three covalent bonds (6 electrons around B)
34
What is an expanded octet (HL)?
When a central atom has more than eight electrons, possible for period 3 elements and beyond.
35
how many valence electrons does diazene (N₂H₂) have?
12 valence electrons.
36
Why does diazene contain an N=N double bond?
Single bonds alone do not give nitrogen a full octet; a double bond is required.
37
How many electrons surround each nitrogen in diazene?N₂H2
- 4 from the N=N double bond
- 2 from the N–H bond
- 2 from one lone pair
Total = 8 (octet satisfied)
38
How does hydrazine (N₂H₄) differ from diazene?
Hydrazine has an N–N single bond, with each nitrogen forming three single bonds and one lone pair.
39
What are limitations of the covalent (Lewis) model?
Lewis structures assume localised electron pairs, so they fail for delocalisation (resonance), give incorrect predicted bond lengths/strengths for resonance systems, cannot explain magnetic behaviour (e.g. O₂ paramagnetism), and do not capture orbital-based bonding and energetics fully.
40
Why are resonance and hybridisation introduced (HL)?
When experimental data (bond length, strength, reactivity) do not match Lewis predictions.
- resonance accounts for delocalised π electrons and hybridisation/orbital overlap explains equivalent bonds and observed geometries.
41
Why is benzene historically important?
Benzene’s equal C–C bond lengths (not alternating) and unusual stability/reactivity could not be explained by localised single/double bonds, driving the development of resonance/delocalised π bonding.
41
exam sentence linking covalent bonding to properties
Covalent bonding stabilises molecules via electrostatic attraction between nuclei and shared electron pairs, while the distribution of electron density and molecular geometry determine polarity and intermolecular forces, which control macroscopic properties such as boiling point, volatility, solubility, and conductivity.
42
One-sentence summary of Lewis structures (perfect for long answers)?
Lewis structures model valence electrons to predict connectivity, formal charge, and electron domains (shape/polarity), while recognising key limitations such as resonance, octet exceptions, and the need for orbital models at HL.
43
What is the difference between single, double and triple covalent bonds?
A single bond involves one shared pair of electrons. A double bond involves two shared pairs of electrons. A triple bond involves three shared pairs of electrons.
44
Why do multiple bonds form in some molecules?
Multiple bonds form when single-bond arrangements cannot satisfy octets (or minimise formal charge), so extra electron pairs are shared to increase bonding electron density between nuclei, lowering energy and stabilising the molecule/ion.
45
Why does O₂ contain a double bond?
Each oxygen atom has six valence electrons; sharing two electron pairs allows both atoms to achieve a full octet.
46
Why does N₂ contain a triple bond?
Nitrogen atoms must share three electron pairs to complete their octets, resulting in a triple bond that makes N₂ very stable.
47
Why is nitrogen gas relatively unreactive despite being abundant in air?
The N≡N bond has very high bond enthalpy (strong σ bond + two π bonds), so breaking it requires a large activation energy; therefore N₂ is kinetically very stable and relatively unreactive.
48
How are multiple bonds represented in Lewis formulas?
One line = single bond. Two lines = double bond. Three lines = triple bond. Each line represents one shared electron pair.
49
Define bond length.
Bond length is the distance between the nuclei of two bonded atoms.
50
Define bond strength.
Bond strength is a measure of how strong a bond is, usually expressed as bond enthalpy, the energy required to break the bond.
51
What is the relationship between bond order and bond length?
Higher bond order → shorter bond length. Triple bonds are shorter than double bonds, which are shorter than single bonds.
52
What is the relationship between bond order and bond strength?
Higher bond order → greater bond strength. Triple bonds are stronger than double bonds, which are stronger than single bonds.
53
Why are multiple bonds shorter and stronger than single bonds?
Multiple bonds involve more shared electrons, increasing electrostatic attraction between the nuclei and the electron cloud, pulling atoms closer together.
54
State the trend in C–C bond length and bond enthalpy from single to triple bonds.
From single → double → triple: Bond length decreases. Bond enthalpy increases.
55
Why is a double bond not twice as strong as a single bond?
Because a double bond consists of one sigma bond and one pi bond, and the pi bond is weaker than a sigma bond.
- it has less effective overlap and electron density further from the bond axis, so it contributes less strength than a σ bond
56
How does bond length change down Group 17 in X₂ molecules?
Bond length increases down the group as atomic radius increases.
57
How does bond enthalpy change down Group 17 in X₂ molecules?
Larger atomic radius increases bond length so the bonding pair is further from both nuclei, reducing electrostatic attraction and orbital overlap, lowering bond enthalpy down the group.
58
Why is F₂ an exception to the halogen bond enthalpy trend?
F atoms are very small, so lone pairs are close together and repel strongly (lone pair–lone pair repulsion), destabilising the bond and reducing F–F bond enthalpy compared with the trend
59
How do double and triple bonds influence reactivity?
Molecules with multiple bonds are generally more reactive because pi bonds are weaker and more easily broken during reactions.
- electron density lies above/below the bond axis, further from the nucleus and is more exposed
60
Why are alkenes and alkynes more reactive than alkanes?
Because they contain double or triple bonds with pi bonds, which are more accessible to reacting species.
61
What is a coordination (dative) bond?
A covalent bond in which both electrons in the shared pair originate from the same atom.
62
How is a coordination bond represented in diagrams?
By an arrow, pointing from the electron-pair donor to the electron-pair acceptor.
63
Give two simple examples of species containing coordination bonds.
NH₄⁺ (N donates a lone pair to H⁺). H₃O⁺ (O donates a lone pair to H⁺).
64
Does a coordination bond differ from other covalent bonds once formed?
No. Once formed, a coordination bond is indistinguishable from other covalent bonds.
65
Describe the bonding in carbon monoxide, CO.
CO contains a triple bond (one σ and two π). The bonding can be described using a coordinate (dative) component in one of the bonding pairs. However, once formed, all bonding pairs contribute to the same electron density and bonds are experimentally equivalent
66
Why are all three bonds in CO considered equivalent?
Because once formed, the coordination bond behaves identically to other covalent bonds.
67
What is a Lewis base?
A species that donates a pair of electrons.
68
What is a Lewis acid?
A species that accepts a pair of electrons.
69
Why do Lewis acid–base reactions form coordination bonds?
Because the Lewis base donates a lone pair to the Lewis acid, forming a shared electron pair.
70
Why can transition metal ions act as Lewis acids?
Transition metal ions can act as Lewis acids because they have empty orbitals in an unoccupied energy level that can accept electron pairs from Lewis bases (ligands). For example, the Fe³⁺ ion can use empty 4s and 4p orbitals to accept electron pairs and form coordination bonds
71
What is a ligand?
A molecule or ion that donates a lone pair to a central metal ion to form a coordination bond.
72
Give examples of common ligands.
H₂O. NH₃. Cl⁻.
73
Name two complex ions containing coordination bonds.
[Fe(H₂O)₆]³⁺. [CuCl₄]²⁻.
74
What does the VSEPR model explain?
It explains and predicts the three-dimensional shape of molecules and ions based on repulsion between electron domains around a central atom.
75
What is the fundamental principle behind the VSEPR model?
Electron domains in the same valence shell repel each other because they contain negative charge, so they arrange themselves as far apart as possible.
76
Why is the term “electron pair” an oversimplification in VSEPR?
Because regions of electron density repel, and a double/triple bond contains more electron density but still acts as a single region; therefore VSEPR uses electron domains (not “pairs”) to predict geometry.
77
Define an electron domain.
An electron domain is any region of electron density around the central atom, including:
a single bond
a double bond
a triple bond
a non-bonding (lone) pair
78
How do you determine the number of electron domains around a central atom?
By examining the Lewis structure and counting each bond (single/double/triple) as one domain and each lone pair as one domain.
79
What determines the electron domain geometry of a molecule?
The total number of electron domains (bonding + non-bonding) around the central atom.
80
What determines the molecular geometry (shape) of a molecule?
The positions of the bonded atoms only, not the lone pairs.
81
Why can electron domain geometry and molecular geometry be different?
Electron-domain geometry includes lone pairs, but molecular geometry describes only atom positions; lone pairs occupy space and repel strongly, altering bond angles while not appearing in the molecular shape.
82
Why do lone pairs cause more repulsion than bonding pairs?
Lone pairs are attracted to only one nucleus, so their electron density is more concentrated near the central atom, increasing repulsion relative to bonding pairs whose density is shared between two nuclei
83
Why do multiple bonds cause more repulsion than single bonds?
Multiple bonds contain higher electron density in the same domain, increasing repulsion compared with a single-bond domain, which distorts bond angles
84
State the order of repulsive strength between electron domains.
lone pair > multiple bond > single bond
85
What effect do lone pairs and multiple bonds have on bond angles?
Lone pairs and multiple bonds increase repulsion, causing deviations from ideal angles: lone pairs compress adjacent bond angles more strongly; multiple bonds can expand some angles and compress others relative to ideal geometry.
86
What geometry results from 2 electron domains?
Linear electron domain geometry.
87
What is the ideal bond angle for 2 electron domains?
180°
88
Give textbook examples of molecules with 2 electron domains.
BeCl₂, CO₂, C₂H₂ (ethyne)
89
What geometry results from 3 electron domains?
Trigonal (triangular) planar electron domain geometry.
90
What is the ideal bond angle for 3 electron domains?
Approximately 120°
91
When is the molecular geometry also trigonal planar?
When all three electron domains are bonding, with no lone pairs (e.g. BH₃, AlCl₃).
92
Why can bond angles in trigonal planar molecules deviate slightly from 120°?
Because unequal domain repulsions occur: a double bond domain repels more than a single bond, and lone pairs repel most, so angles adjust to maximise separation, causing slight deviations from 120°
93
Explain the bond angle distortion in methanal (H₂CO).
The C=O double bond is a higher-density domain, so it repels adjacent C–H domains more strongly, increasing H–C=O angles and compressing H–C–H slightly below 120°.
94
What are the electron domain geometry and molecular geometry of ozone (O₃)?
Electron domain geometry: trigonal planar
Molecular geometry: bent (V-shaped)
95
Why is the O–O–O bond angle in ozone less than 120° (≈117°)?
Ozone has trigonal planar electron geometry, but one lone pair plus a high-density bonding domain increases repulsion asymmetrically; lone-pair repulsion compresses the O–O–O bond angle below 120° (≈117°)
96
What geometry results from 4 electron domains?
Tetrahedral electron domain geometry.
97
What is the ideal bond angle for a tetrahedral arrangement?
109.5°
98
When are electron domain geometry and molecular geometry both tetrahedral?
When all four domains are bonding, e.g. CH₄.
99
What is the molecular geometry of NH₃ and its bond angle?
Molecular geometry: trigonal pyramidal
Bond angle: approximately 107°
100
What is the molecular geometry of H₂O and its bond angle?
Molecular geometry: bent (V-shaped)
Bond angle: approximately 104.5°
101
Why is the bond angle in H₂O smaller than in NH₃?
H₂O has two lone pairs (vs one in NH₃), so total lone-pair repulsion is greater, compressing the H–O–H bond angle more (≈104.5° vs ≈107°)
102
Why must the Lewis structure be drawn before using VSEPR?
Because the number of bonding and non-bonding electron domains can only be determined from the Lewis structure.
103
State the correct IB step-by-step method for predicting molecular shape.
Draw the Lewis structure
Count total electron domains on the central atom
Determine electron domain geometry
Determine molecular geometry from bonding domains
Consider lone pair and multiple bond distortions
104
How useful is the VSEPR model?
It is very useful for predicting basic molecular shapes and approximate bond angles.
105
What are the limitations of the VSEPR model?
VSEPR predicts approximate geometry/angles but cannot quantitatively predict exact bond angles or explain anomalies without additional ideas (e.g. resonance, multiple-bond effects, and orbital/hybridisation models)
106
How do you predict and explain bond angles using the VSEPR model in IB exams?
1. State the ideal bond angle from the electron domain geometry
(linear = 180°, trigonal planar = 120°, tetrahedral = 109.5°).
2. say “slightly less/more” and justify using relative repulsion: lone pair > multiple bond > single bond
3. Justify using electron-domain repulsion, explaining that lone pairs and multiple bonds cause greater repulsion than single bonding pairs, leading to bond-angle distortion. use "slightly less" "slightly more
107
bond polarity
Bond polarity arises from electronegativity difference causing unequal sharing of the bonding pair and creating a bond dipole (δ⁺ → δ⁻), with dipole strength increasing as ΔEN increases.
108
Why are not all covalent bonds equally shared?
Different electronegativities mean one nucleus attracts the bonding pair more strongly, shifting electron density and producing partial charges rather than equal sharing.
109
What is electronegativity?
Electronegativity is a measure of the ability of an atom to attract electrons in a covalent bond.
110
Which scale is used to describe electronegativity?
The Pauling scale, which ranges approximately from 0 to 4.
111
Where are electronegativity values found in the IB exam?
In the IB data booklet.
112
What happens to the shared electrons when one atom is more electronegative?
Electron density shifts toward the more electronegative atom, giving it δ⁻ and the other δ⁺, producing a bond dipole with magnitude related to ΔEN.
113
What is meant by a polar covalent bond?
A polar covalent bond is a covalent bond with an unsymmetrical distribution of electron density.
114
What symbol is used to represent partial charges in polar bonds?
The Greek letter δ (delta).
115
What does a δ charge represent?
A partial charge that has no fixed value and is always less than the unit charge of an ion such as X⁺ or X⁻.
116
Which atom becomes δ⁻ in a polar bond?
The more electronegative atom, because it has the greater share of the electrons.
117
Which atom becomes δ⁺ in a polar bond?
The less electronegative atom, because it has a smaller share of the electrons.
118
What is a bond dipole?
A bond dipole refers to the separation of partial positive and partial negative charges within a polar covalent bond.
119
How can the polar nature of a bond be shown in a structure?
By writing δ⁺ and δ⁻ on the atoms, or by drawing a dipole arrow pointing towards the more electronegative atom.
120
Why is the H–Cl bond polar?
Chlorine is more electronegative than hydrogen, so the shared electrons are pulled closer to chlorine, giving Hδ⁺–Clδ⁻.
121
Why are the O–H bonds in water polar?
Oxygen is more electronegative than hydrogen, so the electron density is greater on the oxygen atom.
122
What determines how polar a covalent bond is?
The electronegativity difference and how much electron density is displaced from the bond midpoint (greater ΔEN → larger dipole moment).
123
State the general trends in electronegativity in the periodic table.
Electronegativity increases across a period; Electronegativity increases up a group.
124
Which element is the most electronegative?
Fluorine, with a value of about 4.0 on the Pauling scale.
125
Arrange the hydrogen halides in order of increasing bond polarity.
H–I < H–Br < H–Cl < H–F.
126
Which covalent bonds are truly non-polar?
Bonds between identical atoms, such as H₂, O₂, and F₂, where the electronegativity difference is zero.
127
What is meant by a “pure covalent” bond?
A bond in which electrons are shared equally, with no bond dipole.
128
Why is the C–H bond often considered non-polar?
Because the electronegativity difference between carbon and hydrogen is very small, so the bond has very low polarity.
129
How does bond polarity relate to ionic character?
Increasing ΔEN increases charge separation, so the bond’s electron density becomes more uneven and the bond shows greater ionic character (approaching complete transfer at the ionic extreme).
130
Why do polar covalent compounds show some properties similar to ionic compounds?
Because the partial separation of charges introduces some ionic character into the covalent bond.
131
Give an example of a polar covalent compound showing ionic-type behaviour.
HCl is polar covalent but ionises in water because the polar H–Cl bond and hydration stabilise ions; HCl(aq) forms H₃O⁺ and Cl⁻, giving mobile ions that conduct electricity.
132
What is meant by the bonding continuum?
Bonding exists on a continuum from pure covalent → polar covalent → ionic, rather than as completely separate types.
133
Why are the boundaries between ionic and covalent bonding described as “fuzzy”?
Because many bonds are partly covalent and partly ionic, so substances are often described as predominantly covalent or predominantly ionic.
134
What is meant by a pure covalent bond?
A pure covalent bond involves equal sharing of electrons (electronegativity difference = 0), so there is no bond dipole.
135
Give examples of pure covalent bonds / molecules.
Bonds between the same atoms, e.g. Cl₂, H₂, O₂, F₂.
136
What is meant by a polar covalent bond?
A polar covalent bond involves unequal sharing of electrons, giving partial charges (δ⁺ and δ⁻) due to a difference in electronegativity.
137
Why is a polar covalent bond described as having partial transfer of electrons?
It’s “partial transfer” because electron density is shifted, not fully transferred; the bond remains covalent but with significant dipole character between the pure covalent and ionic extremes.
138
What is meant by an ionic bond in the bonding continuum diagram?
An ionic bond involves complete transfer of electrons, forming a lattice of oppositely charged ions (e.g. NaCl).
139
State the bonding continuum shown in the textbook.
As electronegativity difference increases: pure covalent → polar covalent → ionic.
140
Why are the boundaries between pure covalent, polar covalent, and ionic described as “fuzzy”?
Because many bonds lie between the extremes, forming intermediate cases, so classification is not always clear-cut.
141
Complete the statement: “In a covalent bond, the greater the difference in electronegativity values…”
“…the more polar the bond will be.”
142
What is meant by “bond dipoles can be shown with partial charges or vectors”?
You can show polarity by writing δ⁺/δ⁻ on atoms or by drawing a dipole arrow showing the direction of electron pull towards the more electronegative atom.
143
molecular polarity
Molecular polarity is the overall polarity of a molecule, determined by the combined effect of its bond polarities and molecular geometry.
144
What two factors determine whether a molecule is polar or non-polar?
- The polar bonds present
- The molecular geometry (orientation of those bonds)
145
Why can a molecule with polar bonds be non-polar overall?
Because if the polar bonds are equal in polarity and symmetrically arranged, their dipoles cancel out. e.g. CO₂, CCl₄).
146
What does it mean when bond dipoles cancel out?
There is no net dipole, so the molecule is non-polar.
147
How can bond dipoles be treated when predicting molecular polarity?
As vectors; the molecular polarity is the resultant of all bond dipoles.
148
Why are CO₂, BF₃, and CCl₄ non-polar molecules?
They contain polar bonds, but their symmetrical geometries cause the dipoles to cancel out.
149
When will a molecule be polar?
When dipoles do not cancel due to asymmetrical geometry and/or different substituents, producing a non-zero resultant dipole moment.
150
What is meant by a net dipole?
A non-zero overall dipole moment that causes a molecule to experience a turning force in an electric field.
151
Why are CH₃Cl, NH₃, and H₂O polar molecules?
Their shapes are not fully symmetrical (and/or bonds differ), so bond dipoles do not cancel: CH₃Cl has a strong C–Cl dipole; NH₃ and H₂O have lone-pair-driven bent/pyramidal shapes giving a net dipole
152
What does it mean for a molecule to be IR active?
It can absorb infrared radiation, causing it to vibrate at a higher frequency.
153
What condition must be met for a molecule to be IR active?
It must have an overall dipole moment associated with bond vibration.
154
Why must diatomic molecules have polar bonds to be IR active?
Because they have only one vibrational mode (stretching), so the bond itself must be polar.
155
Why can polyatomic molecules be IR active even if they contain no polar bonds?
A vibration can create a temporary dipole moment (change in dipole) even if the molecule is nonpolar overall; IR absorption occurs when the vibration causes a change in dipole moment.
156
Why can CH₄ show IR absorption even though it is non-polar overall?
Some stretching and bending vibrations produce a temporary change in dipole moment, allowing IR absorption.
157
What is a covalent network structure?
A crystalline lattice in which atoms are linked together by covalent bonds in a regular repeating pattern, forming a structure with no finite size (effectively one “giant molecule”).
158
What are covalent network structures also commonly called?
Giant molecular or macromolecular structures.
159
How do most covalent substances differ from covalent network substances?
Most covalent substances exist as discrete molecules with a finite number of atoms; covalent network substances form a continuous lattice.
160
Why do covalent network structures have very different properties from small covalent molecules?
Network solids have covalent bonds throughout a giant lattice, so melting/boiling requires breaking many strong covalent bonds (high mp/bp); small molecules are held by intermolecular forces, so they melt/boil at much lower temperatures.
161
Define allotropes.
Allotropes are different bonding and structural patterns of the same element in the same physical state, giving different chemical and physical properties.
162
Give an example of allotropes mentioned (not carbon).
O₂ and O₃ are allotropes of oxygen (both gases).
163
Which carbon allotropes must you describe and compare?
Diamond, graphite, fullerenes, graphene. (and silicon + SiO₂ later in Part 2)
164
Describe the structure of diamond (bonding + geometry).
Each carbon is sp³ hybridized and covalently bonded to 4 other carbons in a tetrahedral arrangement with bond angles of 109.5°, in a regular repetitive pattern.
165
Why does diamond not conduct electricity?
It is a non-conductor of electricity because all electrons are bonded and so non-mobile.
166
What is diamond’s thermal conductivity (relative statement from the table)?
Diamond conducts heat extremely well because its rigid 3D covalent lattice transmits vibrational energy (phonons) efficiently through strong covalent bonds
167
Describe diamond’s appearance
Highly transparent, lustrous crystal.
168
State key physical/chemical properties of diamond .
Hardest known natural substance; cannot be scratched by other materials; brittle; very high melting point.
169
Give uses of diamond
Polished for jewellery and ornamentation; tools and machinery for grinding and cutting glass.
170
Describe the structure of graphite (bonding + layers).
Each carbon is sp² hybridized and covalently bonded to 3 others, forming hexagons with bond angles of 120°, arranged in parallel layers.
171
What happens to the remaining valence electron in graphite?
The remaining valence electron on each carbon is delocalized, so it can move freely across the layer.
172
What holds graphite layers together, and what does this allow?
Layers are held only by weak London dispersion forces, so they can slide over each other.
173
Explain graphite’s electrical conductivity (table wording idea).
Graphite is a good electrical conductor; it contains one non-bonded, delocalized electron per atom that gives electron mobility.
174
Describe graphite’s thermal conductivity (table wording idea).
Graphite conducts heat well along layers (delocalised electrons + strong in-plane bonding) but poorly between layers because weak intermolecular forces limit energy transfer perpendicular to the layers
- not a good conductor unless the heat can be forced to conduct in a direction parallel to the crystal layers
175
Describe graphite’s appearance (table).
Non-lustrous, grey crystalline solid.
176
Explain graphite’s softness/slipperiness (table).
It is soft and slippery due to slippage of layers over each other; brittle; very high melting point; the most stable allotrope of carbon.
177
Give uses of graphite
A dry lubricant; in pencils; electrode rods in electrolysis.
178
Describe graphene’s structure
Each carbon is sp² hybridized and bonded to 3 others forming hexagons with 120° bond angles, but it exists as a single layer (a two-dimensional material), often described as a honeycomb/chicken wire structure.
179
What happens to the remaining valence electron in graphene?
The remaining valence electron on each carbon is delocalized.
180
Describe graphene’s electrical conductivity .
Very good electrical conductor; one delocalized electron per atom gives electron mobility across the layer.
181
Describe graphene’s thermal conductivity
Best thermal conductivity known, even better than diamond.
182
Describe graphene’s appearance.
Almost completely transparent.
183
State key physical properties of graphene
Thickness of just one atom; the thinnest material; also the strongest (about 100× stronger than steel); very flexible; very high melting point.
184
How are graphene, graphite, nanotubes, and fullerenes related?
A single separated layer of graphite is graphene; a rolled-up layer of graphene is a nanotube; a closed cage is a fullerene.
185
graphene’s discovery and current applications?
Although graphene’s conductivity/strength make it promising for composites, sensors, and electronics, large-scale use is limited by difficulty producing large, defect-free sheets and integrating them into devices consistently.
186
Describe the structure of a C₆₀ fullerene (table).
Each carbon is sp² hybridized and bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons; it is a closed spherical cage where each carbon is bonded to 3 others.
187
What important note is made about C₆₀ fullerenes in the table?
It is not a giant molecule (it has a fixed formula).
188
Describe electrical conductivity of C₆₀ fullerenes
C₆₀ fullerenes are poor conductors of electricity; although individual molecules have delocalized electrons, there is little electron movement between molecules.
189
Describe thermal conductivity of C₆₀ fullerenes
Very low thermal conductivity.
190
Describe appearance of C₆₀ fullerenes
Black powder.
191
Give physical/chemical properties and uses of fullerenes
Very light and strong; reacts with potassium to make superconducting crystalline material
- low melting point
- uses include lubricants, medical and industrial devices for binding specific target molecules
- related forms make nanotubes and nanobuds used as capacitors in the electronics industry, and catalysts.
192
Which group is silicon in, and how many valence electrons does it have?
Silicon is a Group 14 element and has four valence electrons.
193
Describe the structure of elemental silicon.
In the elemental state, each silicon atom is covalently bonded to four other silicon atoms in a tetrahedral arrangement, forming a giant covalent lattice structure similar to diamond.
194
How is silicon’s structure similar to diamond?
Both have a tetrahedral network in which each atom forms four covalent bonds, producing a rigid three-dimensional lattice.
195
Why does elemental silicon form a covalent network structure rather than discrete molecules?
Because each silicon atom uses all four valence electrons to form covalent bonds, extending throughout the lattice.
196
What is the electrical conductivity of pure silicon at low temperatures?
Pure silicon is a poor conductor of electricity at low temperatures.
197
Why does the electrical conductivity of silicon increase as temperature increases?
Thermal energy frees some electrons from covalent bonds, allowing them to move through the lattice and conduct electricity.
198
What is meant by a “hole” in silicon’s structure?
When an electron is freed from a covalent bond, it leaves behind a positively charged hole, which other electrons can move into, enabling charge flow.
199
How does this process allow silicon to conduct electricity?
Electron movement and hole movement repeat throughout the lattice, allowing charge carriers to move and conduct electricity.
200
What is meant by “doping” silicon?
Adding very small amounts of Group 13 or Group 15 elements to silicon to increase its electrical conductivity.
201
Give examples of elements used to dope silicon.
Boron (Group 13) and arsenic (Group 15).
202
Why is silicon so important in electronics?
Because its conductivity can be precisely controlled by temperature and doping, making it ideal for semiconductors, computer chips, and transistors.
203
What is silicon dioxide (SiO₂) commonly known as?
Silica or quartz.
204
What type of structure does silicon dioxide form?
A giant covalent (network) structure.
205
Describe the bonding in silicon dioxide.
Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is covalently bonded to two silicon atoms, forming a continuous three-dimensional network.
206
How can the role of oxygen atoms in SiO₂ be described?
Oxygen atoms act as bridges between tetrahedrally bonded silicon atoms.
207
Why is the formula SiO₂ described as an empirical formula?
in a giant covalent lattice there is no discrete “molecule,” so SiO₂ gives only the simplest ratio of atoms in the continuous network; the real structure contains an effectively infinite number of atoms.
208
Why is the SiO₂ structure very strong?
Because all four silicon Each Si forms four strong Si–O covalent bonds in a 3D tetrahedral network, so many bonds must be broken to deform/melt the lattice, giving high strength and high melting point.
209
State the key properties of silicon dioxide (4).
* Strong
* Insoluble in water
* High melting point
* Non-conductor of electricity
210
What everyday materials are different forms of silicon dioxide?
Glass and sand.
211
What key bonding difference exists between carbon and silicon in network structures?
Carbon strongly catenates (stable C–C bonds) forming diverse networks/chains, whereas silicon forms weaker Si–Si bonds and is stabilised in nature mainly as strong Si–O networks (silicates/silica).
212
What is catenation?
The ability of an element (especially carbon) to form covalent bonds with itself, creating chains, rings, or networks.
213
Why does carbon show greater catenation than silicon?
C–C bonds are stronger due to better orbital overlap from carbon’s smaller radius; Si–Si bonds are weaker, so long Si chains are less stable and Si prefers bonding to O.
214
State the bond enthalpy values given for C–C and Si–Si bonds.
* C–C bond enthalpy ≈ 346 kJ mol⁻¹
* Si–Si bond enthalpy ≈ 226 kJ mol⁻¹
215
Why are C–C bonds stronger than Si–Si bonds?
Carbon has a smaller atomic Carbon’s smaller atomic radius allows shorter bond length and more effective overlap, increasing electron density between nuclei and strengthening electrostatic attraction compared to longer Si–Si bonds.
216
How does the strength of the Si–O bond compare to the Si–Si bond?
The Si–O bond is much stronger.
217
How does bond enthalpy of Si-O bond explain the abundance of silicon dioxide in nature?
Because the strong Si–O bond favors the formation of SiO₂ network structures, making silica very stable.
218
Why is silicon called a semiconductor?
Because its electrical conductivity lies between conductors and insulators and can be controlled by temperature and doping.
219
Give a Level 7 summary comparing carbon and silicon network structures.
Carbon forms a wide range of covalent network allotropes due to strong C–C bonding and extensive catenation, while silicon forms similar tetrahedral networks but preferentially bonds to oxygen because Si–O bonds are much stronger than Si–Si bonds, leading to stable silicon dioxide structures with high melting points and low electrical conductivity.
220
expanded octet
An expanded octet occurs when a central atom has more than eight electrons in its valence shell.
221
Which atoms can form expanded octets?
Atoms from Period 3 and beyond (e.g. P, S, Cl, Xe), because they have accessible empty orbitals.
222
Why can Period 2 elements not form expanded octets?
Their valence shell contains only s and p orbitals, so they cannot accommodate more than eight electrons.
223
Why is the octet rule described as a guideline rather than a law?
Because atoms can have incomplete octets, expanded octets, or odd numbers of electrons, depending on stability.
224
What electron-domain geometry results from five electron domains?
Trigonal bipyramidal.
225
What are the ideal bond angles in a trigonal bipyramidal arrangement?
120° between equatorial positions, 90° between axial and equatorial positions, 180° between axial positions.
226
Why are equatorial positions preferred by lone pairs in trigonal bipyramidal structures?
Lone pairs occupy equatorial positions because equatorial sites have fewer 90° interactions; this minimises lone pair–bond pair repulsion and lowers the system’s energy.
227
What is the molecular geometry when there are 5 bonding pairs and 0 lone pairs?
Trigonal bipyramidal (e.g. PCl₅).
228
What molecular geometry results from 4 bonding pairs and 1 lone pair?
Seesaw (unsymmetrical tetrahedron).
229
Why is the seesaw shape distorted from ideal angles?
The lone pair creates stronger repulsion than bonding pairs, especially at 90°, so bond angles deviate from ideal TBP values (ax–eq < 90°, eq–eq < 120°) as bonding domains are pushed away from the lone pair.
230
What molecular geometry results from 3 bonding pairs and 2 lone pairs?
T-shaped.
231
Where are the lone pairs positioned in a T-shaped molecule?
Both lone pairs occupy equatorial positions to minimize repulsion.
232
What molecular geometry results from 2 bonding pairs and 3 lone pairs?
Linear.
- each lone pair repels the otehr equally
-symmetrically around central atom
233
Why is a five-domain molecule with three lone pairs linear?
Three lone pairs occupy the three equatorial positions to minimise 90° repulsions, leaving the two bonding pairs axial and 180° apart, giving a linear molecular geometry.
234
What electron-domain geometry results from six electron domains?
Octahedral.
235
What are the ideal bond angles in an octahedral arrangement?
90° (and 180° between opposite bonds).
236
What molecular geometry results from 6 bonding pairs and 0 lone pairs?
Octahedral (e.g. SF₆).
237
xWhat molecular geometry results from 5 bonding pairs and 1 lone pair?
Square pyramidal.
238
Where is the lone pair positioned in a square pyramidal molecule?
In an axial position, pushing the bonded atoms slightly downward.
239
Why are bond angles in square pyramidal molecules slightly less than 90°?
Because lone-pair repulsion is greater than bonding-pair repulsion.
240
What molecular geometry results from 4 bonding pairs and 2 lone pairs?
Square planar.
241
Where are the two lone pairs located in square planar molecules?
In opposite axial positions, minimizing repulsion.
242
Why does XeF₄ form a square planar shape?
Two lone pairs occupy opposite positions, leaving four bonding pairs in a square plane.
243
How is molecular polarity determined in molecules with five or six electron domains?
By molecular geometry and symmetry, not just bond polarity.
244
Why is PCl₅ non-polar?
Its trigonal bipyramidal geometry is symmetrical, so dipoles cancel.
245
Why is SF₄ polar?
SF₄ has a seesaw molecular geometry due to one lone pair in a trigonal bipyramidal electron arrangement; the S–F bond dipoles are not symmetrically arranged, so they do not cancel, giving a net dipole.
246
Why is XeF₄ non-polar?
Its square planar geometry is symmetrical, causing dipole cancellation.
247
Give a Level-7 sentence explaining expanded octets.
Expanded octets occur in Period 3 and heavier elements because their valence shells can accommodate more than eight electrons, allowing additional electron domains and a wider range of molecular geometries.
248
Give a Level-7 sentence linking lone-pair position to molecular shape.
In molecules with five or six electron domains, lone pairs occupy positions that minimize repulsion, which determines molecular geometry and causes deviations from ideal bond angles.
249
Give a Level-7 sentence linking geometry to polarity.
Molecular polarity in expanded-octet molecules depends on whether the molecular geometry is symmetrical, as only symmetrical arrangements allow bond dipoles to cancel completely.
250
Why are expanded octets possible for Period 3 and heavier atoms (textbook reason)?
Expanded octets are possible for Period 3 and heavier elements because their valence region can accommodate more than eight electrons. These atoms have available orbitals at similar energy levels, allowing additional electron density to be accommodated beyond an octet during bonding.
Period 2 elements cannot expand their octet because their valence shell is limited to a maximum of eight electrons.
251
What does the textbook say limits expanded octets in Period 2?
Period 2 atoms are limited to 2s and 2p valence orbitals, so they cannot accommodate more than 8 valence electrons
252
What does the textbook say about why the octet rule is “special” for most atoms?
The octet is associated with a special stability for most atoms (a full valence shell), but it isn’t universal because of exceptions like expanded octets.
253
What extra “rule” does the textbook give for predicting shapes with non-bonding pairs (beyond just counting domains)?
Predict arrangements by minimising 90° repulsions in this order: minimise lone pair–lone pair 90° interactions first (highest repulsion), then minimise lone pair–bond pair 90° interactions.number of lone pair–bonding pair 90° angles.
254
In trigonal bipyramidal structures, why do lone pairs go equatorial (exam wording)?
Equatorial positions minimise 90° repulsions, so they reduce repulsion more than axial positions.
255
Why are bond angles often less than ideal values in expanded-octet molecules with lone pairs?
Lone pairs have higher electron density and repel bonding domains more strongly, pushing bonds closer together and compressing angles below ideal values.
256
Why are equatorial positions preferred for lone pairs in trigonal bipyramidal structures?
Equatorial positions have fewer 90° interactions, reducing lone pair–lone pair and lone pair–bonding pair repulsions.
257
In expanded-octet molecules, what determines molecular geometry rather than electron domain geometry?
The positions of the bonded atoms only; lone pairs affect angles but are not counted in the shape.
258
What is formal charge used for in the covalent (Lewis) model?
Formal charge is used to compare possible Lewis structures by estimating charge separation; the preferred structure usually has the smallest magnitude formal charges and places negative charge on more electronegative atoms, corresponding to a lower-energy contributor.
259
When do you usually need formal charge?
When you can draw more than one acceptable Lewis formula for the same species (same atom arrangement) and you need to decide which is most stable / preferred.
260
What key idea does formal charge assume about covalent bonding?
It treats bonds as if electrons are shared equally (like “pure covalent”), so it ignores electronegativity differences during the calculation.
261
What are the two assumptions used to assign electrons to an atom in a Lewis formula (for formal charge)?
(a) Each atom gets an equal share of each bonding pair (1 electron per bond, even for a coordinate bond).
(b) Each atom owns all its lone-pair electrons completely.
262
State the formal charge formula .
FC=V−(1/2B +L)
where V = valence electrons, B = bonding electrons, L = lone-pair electrons.
262
formal charge formula
Formal charge = (valence electrons in the unbonded atom) − (number of electrons assigned to the atom in the Lewis formula).
- for the number of valence electrons, each atom is considered seperately, their valence electrons are added seperately
263
How do you calculate “electrons assigned” to an atom
Electrons assigned = ½ × (electrons in bonding pairs) + (electrons in lone pairs).
264
What does a low formal charge generally suggest about a structure?
Lower magnitude formal charges imply less charge separation and usually lower potential energy, so that structure is generally more stable and contributes more to the resonance hybrid.
- less charge transfer has taken place in forming the structure from its atoms
265
What is the most stable / preferred Lewis formula rule-of-thumb
Prefer the structure with:
the lowest formal charges, and
negative formal charges on the more electronegative atoms (if charges must exist)
266
What must the sum of formal charges equal?
The sum of all FC values equals the overall charge of the species (0 for neutral molecules; the ion charge for ions).
267
What’s the difference between equivalent vs non-equivalent Lewis formulas
- Equivalent Lewis formulas: same numbers of single/multiple bonds (e.g. ozone resonance forms).
- Non-equivalent Lewis formulas: different numbers of single/multiple bonds (e.g. SO₂ can be drawn with different bonding patterns).
Formal charge is especially useful to compare non-equivalent possibilities.
268
Why can resonance structures still matter even if one has the lowest formal charge?
The real structure is a resonance hybrid; even if one contributor has the best formal charges, other valid contributors still contribute, affecting bond order, bond length equalisation, and charge delocalisation
269
Why is formal charge a useful tool but “not the full picture”?
FC is a helpful comparison tool because it assumes equal sharing, but real electron density depends on electronegativity and delocalisation; FC predicts best Lewis contributors, not actual partial charges
270
How is formal charge contrasted with oxidation number (from the right page)?
FC assumes equal covalent sharing (splits bonding pairs), while oxidation states assume ionic transfer to the more electronegative atom; both are models used for different purposes and neither equals real charge exactly.
Neither perfectly describes reality, but both are useful depending on context.
271
Example calculation: Find the formal charge on O in OH⁻.
O has V = 6.
In OH⁻, O has 3 lone pairs → L = 6 electrons, and one single bond → B = 2 bonding electrons.
FC(O)=6−(1/2⋅2+6)=6−(1+6)=−1
So O = −1 (H = 0). Sum = −1 overall
272
Why is the XeO₃ structure with three Xe=O double bonds preferred over the one with three single bonds?
The Xe=O double-bond structure is preferred because it minimises formal charges (closest to zero), reducing charge separation and giving the most stable (lowest-energy) Lewis contributor.
273
formal charge to choose between two possible Lewis formulas for N₂O?
Prefer the Lewis structure that places negative formal charge on oxygen (more electronegative), as this corresponds to a more stable contributor for the resonance hybrid.
274
Use formal charge to explain why BF₃ is an exception to the octet rule
if BF₃ is drawn with three single B–F bonds, boron has only 6 electrons (incomplete octet), but the formal charges are all 0:
- B: V=3, B=6 (3 bonds), L=0 → FC = 3 − (3 + 0) = 0
- Each F: V=7, B=2, L=6 → FC = 7 − (1 + 6) = 0
If you force boron to have an octet by drawing B=F double bonds, you introduce unfavourable formal charges (e.g. F becomes +1 and B becomes −1) which is less stable because it creates charge separation and puts positive charge on very electronegative F.
So BF₃ prefers an incomplete octet because it minimises formal charge.
275
What is a sigma (σ) bond?
A sigma bond forms by the head-on overlap of atomic orbitals, where electron density is concentrated along the bond axis between the nuclei.
275
What are the two types of covalent bond
Sigma (σ) bonds and pi (π) bonds.
276
What is meant by the bond axis?
The bond axis is an imaginary line joining the nuclei of the two bonded atoms.
277
Which types of atomic orbitals can overlap to form a sigma bond?
Sigma bonds can form from overlap of s orbitals, p orbitals, or hybrid orbitals, provided the overlap is head-on along the bond axis.
278
What key feature of electron density characterises a sigma bond?
In a sigma bond, electron density is concentrated directly between the nuclei along the bond axis.
279
What is a pi (π) bond?
A pi bond forms by the lateral (sideways-on) overlap of p orbitals, producing electron density on opposite sides of the bond axis.
279
What type of bond is always present in a single covalent bond?
A single covalent bond always consists of one sigma bond.
280
Why is a sigma bond always formed first in covalent bonding?
Because head-on overlap along the bond axis gives the greatest overlap of orbitals and the strongest attraction between nuclei and shared electrons.
- it gives the highest electron density between nuclei and therefore strongest stabilising nucleus electron attraction
281
Which orbitals are required to form a pi bond?
Pi bonds form only from the lateral overlap of p orbitals.
282
How many sigma and pi bonds are present in a double covalent bond?
A double bond contains one sigma bond and one pi bond.
282
Where is the electron density located in a pi bond?
Electron density is concentrated in two regions above and below the bond axis, not directly between the nuclei.
283
How many sigma and pi bonds are present in a triple covalent bond?
A triple bond contains one sigma bond and two pi bonds.
283
Why can a pi bond not form on its own?
A π bond requires parallel p orbitals on two atoms to overlap sideways, which is only possible when the atoms are already held at the correct distance and orientation by a σ bond framework
283
In which types of covalent bonds do pi bonds occur?
Pi bonds occur only in double bonds and triple bonds.
284
Why are pi bonds weaker than sigma bonds?
Sideways overlap is less effective, placing electron density further from the nuclei and reducing stabilising attraction, so π bonds have lower bond enthalpy and break more easily.
284
What type of overlap forms the H–H bond in H₂?
Head-on overlap of two s orbitals, forming a sigma bond.
285
How does pi-bond weakness affect chemical reactivity?
Because π electrons are more exposed and less strongly held, they are attacked/broken first during reactions, making multiple bonds more reactive than σ-only single bonds.
286
How does the presence of a pi bond influence carbon–carbon reactivity?
Carbon–carbon double bonds are more reactive than single bonds because the pi bond breaks more easily.
286
What type of overlap forms the H–Cl bond in HCl?
Head-on overlap of an s orbital and a p orbital, forming a sigma bond.
287
What type of overlap forms the Cl–Cl bond in Cl₂?
Head-on overlap of two p orbitals, forming a sigma bond.
288
What type of overlap forms C–H bonds in CH₄?
Head-on overlap of a hybrid orbital and an s orbital, forming sigma bonds.
289
What bonds are present in the C≡C bond in ethyne (C₂H₂)?
One sigma bond (from head-on overlap) and two pi bonds (from sideways overlap of p orbitals).
290
What type of overlap forms the sigma bond in C=C in ethene (C₂H₄)?
Head-on overlap of two sp² hybrid orbitals (sp²–sp²) to form the σ bond. The π bond is from sideways p–p overlap.
291
Give an inorganic example of a molecule containing a pi bond from the page.
One of the C=O bonds in the carbonate ion, CO₃²⁻, contains a pi bond.
292
Why is orbital overlap important for explaining molecular shape?
Orbital overlap requires specific orientations in 3D, so hybrid orbitals directed in space explain observed bond angles (e.g. 109.5°, 120°, 180°) that spherical s orbitals alone cannot account for
293
Why does the sigma–pi model go beyond the VSEPR model?
VSEPR predicts shape, but sigma–pi bonding explains how orbitals overlap to produce those shapes and bond strengths.
- σ/π and hybridisation explain why those geometries occur by describing orbital orientation and overlap, and they account for bond strength and equivalence (e.g. identical C–H bonds in CH₄)
294
What is hybridisation?
Hybridisation is the mixing of atomic orbitals on the same atom (s and p at IB level) to form degenerate hybrid orbitals oriented to maximise σ overlap, producing stronger bonds and explaining molecular geometry
- energy of hybrid orbitals is intermediate between the ones mixed
294
Why is hybridisation required to explain bonding in carbon?
Carbon’s ground state has only two unpaired electrons, but carbon forms four equivalent σ bonds; promotion (excitation) creates four unpaired electrons and hybridisation forms four equivalent orbitals, matching observed identical bond lengths/angles (e.g. CH₄) Hybridisation explains how carbon forms four equivalent bonds by first undergoing excitation and then orbital mixing.
295
What must you be able to predict using hybridisation?
The geometry around an atom from its hybridisation, and vice versa.
295
types of hybridisation
sp, sp², and sp³ hybridisation.
296
Why is carbon important in the study of hybridisation?
Because carbon forms a vast number of covalently bonded compounds and forms four covalent bonds in all of them.
297
What is the electron configuration of carbon in its ground state?
1s² 2s² 2p²
(with only two singly occupied orbitals available for bonding)
298
Why does the ground-state electron configuration of carbon fail to explain four bonds?
Because it predicts only two singly occupied orbitals, yet carbon forms four covalent bonds.
299
What does the formation of four covalent bonds by carbon indicate about its electron configuration?
That carbon’s ground-state electron configuration changes during bonding.
300
What is excitation in the context of covalent bonding?
Excitation is the process in which an electron is promoted from the 2s orbital to the vacant 2p orbital within the atom.
301
What is the result of excitation in carbon?
The atom now has four singly occupied orbitals available for bonding.
302
Which rule explains the electron arrangement after excitation?
Hund’s rule, which states that electrons in singly occupied orbitals within the same subshell have parallel spin.
303
Why is excitation energetically allowed during bond formation in carbon?
Excitation costs energy, but during bond formation the energy released by forming additional strong σ bonds (greater overlap) more than compensates, so the overall process is energetically favourable.
304
when does excitation occur. give example
when an electron is promoted from a lower energy atomic orbital to a higher energy atomic orbital within the same energy level
ground state: 1s2, 2s2, 2p2
excited state: 1s2, 2s1, 2p3
305
Why does excitation alone not explain carbon forming four identical bonds?
Excitation gives one 2s and three 2p orbitals of different energies, so bonds would be unequal; hybridisation mixes them into equivalent orbitals, explaining identical bond enthalpies and lengths
306
What observation in methane supports the idea of hybridization?
All four C–H bonds in methane are identical in length and strength.
307
What happens to atomic orbitals during hybridization?
One s and some p orbitals combine to form degenerate hybrid orbitals with new orientations that maximise overlap; any unused p orbitals remain unhybridised and can form π bonds.
308
How are hybrid orbitals different from the original atomic orbitals?
Hybrid orbitals have different energies, shapes, and orientations in space compared to the original parent orbitals
309
Why do hybrid orbitals form stronger covalent bonds?
They allow greater orbital overlap, increasing electron density between the nuclei.
- Hybrid orbitals have directed lobes and greater overlap with partner orbitals, increasing electron density between nuclei and strengthening nucleus–electron attraction (higher bond enthalpy)
310
Which atomic orbitals are involved in hybridization at ib level
s and p orbitals only.
311
How does hybridization explain why bonds in methane are identical?
Hybridization produces four equivalent hybrid orbitals, so all four bonds formed are identical.
312
What does sp³ hybridization involve?
The mixing of one s orbital and three p orbitals to form four identical sp³ hybrid orbitals.
313
What does sp² hybridization involve?
The mixing of one s orbital and two p orbitals to form three identical sp² hybrid orbitals.
314
What does sp hybridization involve?
The mixing of one s orbital and one p orbital to form two identical sp hybrid orbitals.
315
What type of bond forms when a hybrid orbital overlaps with another orbital?
A sigma (σ) bond.
316
What determines the shape of a molecule according to hybridization?
The orientation of the hybrid orbitals in space.
317
Why is hybridisation needed to explain bonding in carbon compounds?
It explains equivalent bonds and correct geometry: e.g. sp³ gives four identical σ bonds in CH₄ (tetrahedral 109.5°), sp² gives trigonal planar σ framework plus one π bond, sp gives linear σ framework plus two π bonds
318
What role does excitation play in hybridisation?
Excitation promotes an electron to create unpaired electrons, allowing the atom to form the required number of covalent bonds before orbital mixing occurs.
319
Describe the excitation of a carbon atom before hybridisation.
An electron is promoted from the 2s orbital to a vacant 2p orbital, producing four singly occupied orbitals available for bonding.
320
Why is excitation energetically favourable in carbon?
Because the energy required for excitation is more than compensated by the energy released when additional strong covalent bonds form.
321
What does “degenerate orbitals” mean in hybridisation?
Degenerate orbitals have the same energy.
322
What is meant by “intermediate energy” hybrid orbitals?
Their energy lies between the energies of the original orbitals from which they were formed.
323
When does sp³ hybridisation occur?
When an atom forms four single (σ) bonds.
324
Which orbitals mix to form sp³ hybrid orbitals?
One s orbital and three p orbitals.
324
How many sp³ hybrid orbitals are formed?
Four equivalent sp³ hybrid orbitals.
325
What is the geometry of sp³ hybrid orbitals?
Tetrahedral.
326
What is the bond angle in sp³ hybridisation?
109.5°.
327
What type of bonds do sp³ hybrid orbitals form?
Sigma (σ) bonds formed by head-on overlap.
- each hybrid orbital overlaps with the atomic orbital of another atom forming four sigma bonds
328
Give an example of a molecule with sp³ hybridisation.
Methane, CH₄.
329
When does sp² hybridisation occur?
When an atom forms a double bond.
330
Which orbitals mix to form sp² hybrid orbitals?
One s orbital and two p orbitals.
331
How many sp² hybrid orbitals are formed?
Three equivalent sp² hybrid orbitals.
332
What is the geometry of sp² hybrid orbitals?
Trigonal planar.
333
What is the bond angle in sp² hybridisation?
120°.
334
What happens to the remaining p orbital in sp² hybridisation?
One unhybridised p orbital remains perpendicular to the plane of the sp² orbitals.
335
How is a π bond formed in sp² hybridisation?
By sideways overlap of unhybridised p orbitals above and below the plane of the molecule.
336
Which bonds are σ and which are π in a C=C double bond?
One σ bond (from sp²–sp² overlap) and one π bond (from p–p sideways overlap).
337
Give an example of a molecule with sp² hybridisation.
Ethene, C₂H₄.
338
When does sp hybridisation occur?
When an atom forms a triple bond.
339
Which orbitals mix to form sp hybrid orbitals?
One s orbital and one p orbital.
340
How many sp hybrid orbitals are formed?
Two equivalent sp hybrid orbitals.
341
What is the geometry of sp hybrid orbitals?
Linear.
342
What is the bond angle in sp hybridisation?
180°.
343
How many unhybridised p orbitals remain in sp hybridisation?
Two unhybridised p orbitals.
344
How are π bonds formed in sp hybridisation?
By sideways overlap of two pairs of unhybridised p orbitals, forming two π bonds.
345
Describe the bonding in a C≡C triple bond.
One σ bond formed by sp–sp overlap and two π bonds formed by sideways p–p overlap.
346
Give an example of a molecule with sp hybridisation
Ethyne, C₂H₂.
347
How is hybridisation related to molecular shape?
Hybrid orbitals orient in space to minimise repulsion and maximise overlap: sp (180° linear), sp² (120° trigonal planar), sp³ (109.5° tetrahedral), which determines observed molecular geometry
348
How does VSEPR link to hybridisation?
Electron-domain count predicts both: 2 domains → sp, 3 → sp², 4 → sp³; VSEPR gives geometry while hybridisation explains orbital orientations that produce it
348
Why do hybrid orbitals form stronger bonds than unhybridised orbitals?
Greater overlap from directional hybrid orbitals concentrates electron density between nuclei, increasing electrostatic attraction and bond enthalpy compared with less directional unhybridised orbitals.
349
Give the mapping between geometry and hybridisation.
Linear → sp
Trigonal planar → sp²
Tetrahedral → sp³
350
Why are π bonds weaker than σ bonds?
Because sideways overlap is less effective than head-on overlap, placing electron density further from the nuclei.
351
Why are molecules with π bonds more reactive?
π electrons are more exposed above and below the bond axis and are easily attacked by electrophiles reactants (those that are attracted to electron dense regions)
-making unsaturated compounds more reactive.
352
What is meant by an unhybridised p orbital?
An unhybridised p orbital is a p orbital that does not mix with s or other p orbitals during hybridisation and retains its original shape and orientation.
353
In sp² hybridisation, how many p orbitals remain unhybridised?
One p orbital remains unhybridised.
354
Why does one p orbital remain unhybridised in sp² hybridisation?
because only one s + two p orbitals mix to form three sp² orbitals. The third p orbital is not used, so it remains a pure p orbital.
- the remaining p orbital stays unhybridised perpendicular to the plane so it can overlap sideways to form the π bond
355
What is the orientation of the unhybridised p orbital in sp² hybridisation?
The unhybridised p orbital is oriented perpendicular (at 90°) to the plane of the sp² hybrid orbitals.
-The three sp² orbitals lie in one flat plane (120° apart). The remaining p orbital points at 90° to that plane, so its lobes stick above and below the plane.
356
Why is the unhybridised p orbital perpendicular to the sp² orbital plane?
Because the three sp² hybrid orbitals lie in one plane, while the remaining p orbital retains its original orientation, which is perpendicular to that plane.
357
importance of the unhybridised p orbital
The unhybridised p orbital is essential for π-bond formation because its perpendicular orientation allows sideways overlap, producing electron density above and below the σ-bond framework and increasing bond strength.
358
What happens to atomic orbitals during hybridisation?
s and p orbitals on the same atom mix to form degenerate hybrids directed for maximum σ overlap; unhybridised p orbitals (if any) remain available for π bonding
359
How do lone pairs affect molecular shape but not hybridisation type?
Hybridisation depends on total electron domains, so lone pairs do not change hybrid type; however, lone pairs repel more strongly and compress bond angles, changing molecular geometry (e.g. sp³ electron geometry tetrahedral, but NH₃ trigonal pyramidal).
360
Do lone pairs ever occupy unhybridised p orbitals in basic hybridisation models?
No. In the hybridisation model used at this level, lone pairs are treated as occupying hybrid orbitals, not unhybridised p orbitals.
361
Explain the role of lone pairs in hybridisation and molecular geometry.
Lone pairs occupy hybrid orbitals and count as electron domains, so hybridisation depends on the total number of domains (bonding + lone pairs) and is unchanged by whether a domain is bonding or non-bonding. However, lone pairs have greater electron density and repel more strongly, compressing bond angles and changing molecular geometry (e.g. tetrahedral electron geometry, trigonal pyramidal shape in NH₃).
IB chemistry HL
flashcards
Decks in class (23)
# Cards
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Atomic structure21
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The periodic table: Classification of Elements145
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2.2-2.3 The Nuclear atom and electronic configurations118
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functional groups organic chemistry1
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the metallic model39
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Particulate nature of matter and separation techniques (Pre-IB)50
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Kinetic molecular theory and changes of state (pre-IB)52
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1 1.4 Counting particles by mass: mole49
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1 1.5 Ideal gases55
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Nuclear atom110
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Electronic configuration123
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the ionic model141
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metallic model43
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polyatomic ions17
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periodic trends80
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The covalent model378
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covalent model pt.2 IMF104
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covalent model - resonance structures + benzene37
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Measuring enthalpy changes18
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1.1121
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1.2167
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..8
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Reaction rates0
