Structure, Bonding and Equilibria Flashcards
Covalent bonding
The formal sharing of electrons between two atoms
- Found in elements with high electronegativity and in heteroatomic molecules with a small difference in electronegativity
Ionic bonding
In compounds where constituent elements have a high difference in elelctronegativity
Metallic bonding
Found in elements with low electronegtaivity (bottom left of periodic table) electrons delocalised into a “sea of electrons”
Hydrogen bonding
Primarily electrostatic, traditionally between a H atom attached to an electronegative atom (non-covalent interaction)
Define elctronegtivity
ability of an atom to attract an electron to itself (Pauling)
Trend of electrongativity down the periodic table
decreases down the table (highest in top right, lowest in bottom left)
Define electropositive
not very electronegative
ΔH of atomisation
The amount of enthalpy change when a compound’s bonds are broken and the component atoms are separated into single atoms
ΔH of reaction
The enthalpy change that accompanies a chemical reaction
1st ionisation enthalpy
The enthalpy change associated with the removal of the first electron from an isolated gaseous atom
2nd ionisation enthalpy
The enthalpy change associated with the removal of the second electron from an isolated gaseous ion
Electron gain enthalpy
The amount of energy released when an electron is added to an isolated gaseous atom
Lattice enthalpy
A measure of the strength of the forces between the ions in an ionic solid
In terms of the van Arkel-Katelaar triangle, what is an ionic bond?
A high diff. in electrongativity
In terms of the van Arkel-Katelaar triangle, what is an covalent bond?
A high electronegativity and a low difference between the two constituents
In terms of the van Arkel-Katelaar triangle, what is an metallic bond?
A low electrongativity with a low difference between constituents
Describe the trends in 1st ionization energy across period 2
- Moving left to right protons increase (nuclear charge increases)
- The outer electrons are all in the same area of space so do not shield each other well (Zeff increases)
- Attraction between nucleus and outer e- increase
- IE increases
Explain the dip in 1st IE at B
New e- in p-orbital does no penetrate 2s very well so increase in Zeff is not as much as expected
Explain the dip in 1st IE at O
Two electrons in same p-orbital, e- repulsion (less energy to remove)
Describe and explain the trend in 1st IE as you move down the periodic table
- IE decreases
- Outer e- further from nucleus
- Less energy required to remove outer e-
- NOT AN EVEN TREND - d-orbitals being filled have a small effect on Zeff due to the 10 e- not effectively shielding the 10 new protons (Zeff increases)
Describe and explain the trend in IE beyond 1st as you move down the periodic table
- 2nd and higher IE is always higher (e- attraction greater to + cation)
- Significant increase in IE when removing from a lower shell (closer to nucleus)
Describe trends in lattice enthalpy
- Lattice enthalpy decreases as ion size increases
- Greater charge on either ion leads to greater lattice enthalpy
- Gain in lattice enthalpy is normally more than economical for amount of energy entering system in IE
- Lattice enthalpy decreases down a group
Describe and explain the trend in covalent bond dissociation energies
As atoms get larger their orbitals increase in size and become more diffuse (BDE decreases) therefore, weaker and longer bond
- Bonds to electronegative atoms are stronger and shorter
Why do covalent bonds form?
- both e- attracted to both positive nuclei
- 2 e- repel eachother
- 2 nuclei repel eachother
- equilibrium in terms of internucleur distance and energy found
What is hybridization?
the combining of s and p orbitals to form sp orbitals (which point to the four corners of the tetrahedron)
What hybridization occurs in an alkane?
sp3 (1s and 3p orbitals) (4 sigma bonds)
What hybridization occurs in an alkene?
sp2 (1s and 2p orbitals) (final p orbital forms pi bond) (3 sigma bonds formed by sp orbitals)
What hybridization occurs in an alkyne?
sp (1s and 1p orbital) (2 sigma bonds forrmed by sp orbitals) (2p orbitals form the two pi bonds at 90 degrees to eachother)
What does an increase in s character in a given bond do to bond length?
Shortens it as s-electrons held closer to the nucleus
Rules for using VSEPR to determine shape
1) Determine which atom you’re determining geometry around
2) Take the group number for this atom and add electrons for substituents (1 for single bond,2 for double etc.)
3) Add/remove e- for charge
4) Divide by 2 to determine no. PAIRS
5) Deduce no. lone pairs and bonding pairs
6) Remember to consider any double or triple bonds before determining the geometry
How does EPG affect hybridization?
Linear - sp
Triagnol planar - sp2
Tetrahedral - sp3
Trigonal bipyramid - sp3d
Octahedral - sp3d2
Steps for drawing a resonance form
1) draw a Lewis dot structure with charges
2) Identify which groups can donate pi e-/lone pairs and those that can accept them
3) Draw curly arrows between the two groups (must be conjugated)
4) ensure nothing has too many bonds
5) draw the new structure
6) Check stability
Define Bronsted-Lowry acid
An acid is a substance which is a proton donor
Define Bronsted-Lowry base
A base is a substance that accepts protons
Define Lewis acid
e- pair acceptor
Define Lewis base
e- pair donor
The formula for the equilibrium constant
Kc = [A-][H3O+]/[HA][H2O]
The formula for acid dissociation constant
Ka = [H+][A-]/[HA]
What does a high Ka indicate in terms of acid strength
High Ka = high acid strength
pKa formula
pKa = -log(Ka)
What does a high pKa indicate in terms of acid strength
low acid strength
Factors affecting acidity
1) strength of H-A bond
2) Electronegativity of A
3) stabilization of A-
4) Solvent
What is an indication of a high stability of A-
- A large no. resonance forms (- charge more dispersed across mol) less likely to reform acid
- If the resonance forms place - on an electronegative atom, increases the stability of the conjugate base
Kb formula
[BH+][OH-]/[B]
pKb formula
pKb = -log(Kb)
Ka of the conjugate base formula
[B][H3O+]/[BH+]
What does a high Ka of the conjugate base indicate about the base strength?
Weak base
What does a high pKa of the conjugate base indicate about the base strength?
Strong base
What does solvent leveling mean?
The effect of strong acids appearing to be equally strong because they all fully dissociate in a good base
Example of a worse base than water to test differences in acidity between strong acids
Ethanoic acid CH3COOH
What will a base stronger than water do in water?
It will deprotonate water to OH- (water acts as an acid)
Kc formula describing self-ionisation of water
Kc = [H3O+][OH-]/[H2O]
define amphiprotic
A substance that can act as both an acid and a base
Self ionisation constant
Kw
Kw expression for water
Kw = [H3O+][OH-] (units: mol2dm-6)
prove water’s pH is 7
at 298K Kw = 10^-14mol2dm-6
In pure water [H3O+]=[OH-]
Therefore [H3O+] = 10^-7moldm-3
pH=-log[H3O+]
pH=-log(10^-7)
pH=7
How does pH of water vary with temperature?
- H2O dissociation is endothermic
- As temp increases pH decreases
- This is always neutral as pure water is always considered “neutral”
pH formula
pH=-log[H3O+]
What defines a strong acid in terms of pKa
pKa<0
How to determine pH of a strong acid (pKa<0)
Because acid is completely dissociated, you can assume that [H3O+(aq)]=[HA]
pH = -log[H3O+]
pH = -log[HA]
Define buffer
Solutions able to resist changes in pH on the addition of an acid or base or on dilution
What is a buffer made up of
A buffer is made up of weak acids and their conjugate bases or weak bases and their conjugate acids
e.g. CH3COOH + CH3COONa, NH3 + NH4Cl
What is the Henderson-Hasslebach equation used for
to find the pH of a buffer
What is the Henderson-Hasslebach equation?
-log(Ka) = -log(pKa) + log([A-]/[HA])
Henderson hasslebach equation
pH = pKa + log([A-]/[HA])
