Structure, Bonding and Equilibria Flashcards
Electronegativity
The ability of an atom to attract an electron to itself.
Electronegative elements
F
O
Cl
Top right of periodic table
Electropositive elements
Cs
Ba
Rb
Bottom left of periodic table
Ionic bonding when…
Difference in electronegativity >2
Covalent bonding when …
0.5< difference in electronegtivity <2
Metallic bonding when…
Difference in electrnegativity <0.5
ΔH atomisation
X(standard) -> X(g)
always endothermic
1st ionisation enthalpy
X(g) -> X⁺(g) + e⁻
always endothermic
Electron gain enthalpy
X(g) + e⁻ -> X⁻(g)
usually exothermic (or close to it)
Lattice enthalpy
X⁺(g) + Y⁻(g) -> XY(s)
always exothermic (when defined in this way)
General trend in I.E. across a period.
Why?
Overall, ionisation energy increases across a period.
Increasing nuclear charge and reduced electron shielding ability increases effective nuclear charge.
Electrons experience greater attraction to nucleus, so more energy input required to remove them.
Zeff
Effective nuclear charge - the actual amount of positive charge experienced by an electron.
Zeff = Z (atomic number) - S (shielding constant)
Electron shielding
Partial neutralisaiton of nuclear charge by core electrons + electron-electron repulsion reduces Zeff felt by outer shell electrons.
Electron penetration
An electron’s relative electron density near the nucleus - i.e. proximity to which an electron can approach the nucleus.
Penetrating power depends on shell and subshell
1s>2s>2p>3s>3p>4s>3d etc
In the second period, why is there a dip in 1st I.E at B and O?
Boron - the outer 2p electron does not penetrate the 2s electrons effectively. Therefore easier to remove.
Nitrogen - repulsion between paired 2p electrons makes them easier to remove.
General trend in I.E down a group.
Why?
Overall, I.E decreases down the group.
Principle quantum number (number of shells) increases. Valence electrons spend more time futher from nucleus - attraction to nucleus is lower - electrons are easier to remove.
In group 1, why is there an uneven trend in I.E between K and Rb?
Between K and Rb the 3d orbitals are filled.
D orbital electrons do not shield the nuclear charge as effectively. I.E for Rb is higher than would be expected as Zeff for valance electrons is higher.
note - I.E for Rb is still lower than for K. However, the difference is smaller than would be expected.
General trend in electron gain enthalpy across a period.
Why?
Overall, electron gain enthalpy becomes more negative across the period.
Zeff increases across the period as nuclear charge increases and electron shielding is less effective.
Makes it easier for atom to accept electron as incoming electron experiences greater attraction.
In period 2, why is there an increase in electron gain enthalpy at N?
The incoming electron has to pair up in a 2p orbital. Must overcome repulsion with paired electron.
Lattice enthalpy is proportional to…
∣z⁺∣∣z⁻∣/(r⁺+r⁻)
absolute product of charge on the anion and cation / sum of radius of anion and cation
Lattice enthalpy increases when…
…ion size decreases
… charge on either ion increases
Bond energy is…
…the energy required to break a specific bond via homolytic fission
units: kjmol-1
always positive
Bond dissociation energies
Standard enthalpy change for the reaction in which the bond in question is broken.
*Average * values
Bonds to larger atoms tend to be…
…weaker and longer.
Bonds to more electronegative atoms tend to be…
…shorter
Equilibrium bond length
Internuclear distance that minimises repulsion between nuclei, but maximises attraction between nuclei and electrons.
Hypervalent species
Species with >8 valence electrons.
Found further down the periodic table.
To which elements is the octet rule not applicable?
Transition metals.
Formal charge
= (e⁻ in valence shell) - (no. of bonds to atoms) - (no. of unshared e⁻)
