Structure and Bonding Flashcards
What are the four main types of bonding?
Ionic, Covalent, Metallic, and Hydrogen bonding.
What is ionic bonding? Give some examples
The bonding occurs in compounds in which the constituent elements have a relatively large difference in electronegativity
Examples are NaCl and CsF
What is covalent bonding? Give some examples
Involves the formal sharing of electrons between two atoms
They are found in elements with high electronegativity e.g. Cl₂, F₂, Br₂
When are covalent bond formed in heteroatomic molecules?
Occurs when the different atoms have a small difference in electronegativity
What is metallic bonding?
Found in elements that have a low electronegativity, with the electrons delocalised into a ‘sea of electrons’
What is hydrogen bonding? Give an example of a molecule where hydrogen bonding occurs
The primary electrostatic interaction between a hydrogen atom attached to an electronegative atom and a second electronegative atom
It is an example of a non-covalent interaction
Water is the main example
What is electronegativity?
The ability of an atom to attract an electron to itself
It is considered a slightly artificial concept as there is no way to directly measure it
What are two common scales that electronegativity is measured on?
The Pauling scale and the Allen scale
Where in the periodic table is electronegativity at its highest?
The top right (F)
Where in the periodic table is electronegativity at its lowest?
The bottom left (Cs)
What is the general trend in electronegativity as you go down the periodic table?
It decreases as the atoms get larger and so are further away from the nucleus
What does ‘electropositive’ mean?
An element with a low electronegativity (the inverse of electronegative)
What does the van Arkel-Ketelaar triangle allow us to predict?
The dominant type of bonding in a compound based on electronegativity and the difference in electronegativity
What type of bonding is present if the electronegativity difference is greater than 2?
Ionic bonding
What type of bonding is present if the electronegativity difference is greater than 0.5 but less than 2?
Polar covalent bonding
What type of bonding is present if the electronegativity difference is less than 0.5?
Covalent bonding
Why do bonds form?
To lower the energy of a system
What does an exothermic reaction tell you about the energy level of the reactants and the products?
The bonds in the product are stronger than those in the starting material - it is said to be thermodynamically favourable
Why is NaCl₂ only a hypothetical compound?
The ionic bonds are not strong enough to compensate for the very high second ionisation enthalpy of Na
The bonds are not strong enough to ‘repay’ the energy needed
What is the general trend of the ionisation energy across a period?
It generally increases
Why does the ionisation energy increase across a period?
The number of protons increases, increasing the nuclear charge
The additional electrons added are in the same area of space (have the same quantum principle number) so do not shield each other as effectively
Why is there a dip in ionisation energy at Boron?
The 2p orbital does not penetrate the 2s orbital very well so the effective nuclear charge is lower
It is not as close to the nucleus leading to less attraction
Why is there a dip in ionisation energy at Oxygen?
There are two electrons in the same p orbital and so there is electron-electron repulsion
Which element has the maximum ionisation energy in period 2, and why does this make it relatively inert?
Neon has the highest ionisation energy, and it cannot form bonds strong enough to compensate for this, making it very hard to react with
What is the trend in ionisation energy as you go down the periodic table?
It decreases
As the principal quantum number increases, electrons spend time further away from the nucleus, decreasing the force of attraction
Why is there a smaller difference in ionisation energy between K and Rb compared to Li to Na?
Electrons have been added in the d orbital between K and Rb
This is due to the 10 electrons d orbital not effectively shielding the 10 added protons, therefore effective nuclear charge is higher
Do successive ionisation energies require more or less energy?
It always takes more energy - this is because it is harder to remove an electron from a cation due to electrostatic forces of attraction
Why is there a much higher increase in the 3rd ionisation energy of Be compared to the 2nd ionisation energy?
There are more protons than electrons, and the electrons are much closer to the nucleus as they have a higher principal quantum number
Why is it impossible for reactions involving ionisation energies of core electrons to occur?
The ionisation energy is so high that it is not possible to recoup the energy for bond formation
Usually, what type of reaction is electron gain enthalpy? Why?
Exothermic (or very close to it)
There is normally sufficient attraction between the nucleus and the ‘new’ electron to overcome electron-electron repulsion
What is the general trend in electron gain enthalpy across a period?
It generally decreases
Why is there an increase in electron gain enthalpy at Be?
The ‘new’ electron is being added into the 2p¹ orbital, and as this is further away from the nucleus, more energy is needed to accept this electron
Why is there an increase in electron gain enthalpy at N?
The ‘new’ electron is being added in the 2p orbital (which is the first that is doubly filled), leading to more electron-electron repulsion, meaning more energy is required to add the electron
What is lattice enthalpy directly proportional to?
( |z⁺| x |z-| ) / (r₊+ r₋)
|z⁺| = Charge on cation
|z-| = Charge on anion
r₊ = Radius of cation
r₋ = Radius of anion
| = Modulus (absolute value)
What is the trend in lattice enthalpy when the anion or cation is larger?
It decreases (lattice enthalpy gets smaller)
What is the trend in lattice enthalpy when the ions have a greater charge? (e.g. Al³⁺ compared to Li⁺)
It increases as the bonds are stronger
However, more energy is needed to complete multiple, successive ionisations
Why are there no known stable compounds of Helium?
The ionisation energy is too high and the making of bonds cannot repay this debt
Why can Xenon and Krypton form stable compounds?
Their ionisation energies are low enough so that the bonds created can repay the debt
Why is the electron gain enthalpy for Noble Gases positive?
The extra electron resides in the next shell much further away from the nucleus, meaning that energy is needed to form the attraction between the nucleus and the outer electron
Why are half-filled shells (p⁴ to p³) more stable than shells that are not half-filled?
Going from p⁴ to p³, there are two electrons in one orbital, and this repulsion means that less energy is needed to form bonds
Why are half-filled shells (p² to p³) more stable than shells that are not half-filled?
It is gaining an electron with all other electrons having the same spin, imparting stability through ‘exchange energy’