Periodic trends

Trends on the periodic table are as follows.

1. Atomic radii: (half the distance between adjacent atoms nuclei)
Decreases across the table and down the groups.
Across the table: electron shells remain the same, but number of protons is increasing = protons have a greater pull of outer electrons.
Down the group: electron shells are increasing so there is electron shielding from the pull of the protons.

2. Ionic radii: (half the distance between adjacent ions nuclei)
Cations are smaller than their atoms as they have lost electrons.
Anions are larger than their atoms as they have gained electrons.

3. Electronegativity: (A measure of the ability of an atom to attract towards itself an electron stared in a chemical bond)
Increases across the period and decreases down the group.
Across the table: Across the table: electron shells remain the same, but number of protons is increasing = protons have a greater pull on bonded electrons
Down the group: electron shells are increasing so there is electron shielding from the pull of the protons.

4. First ionization energy: (energy needed to remove the least tightly held electron form an atom in gaseous state)
Increases across the period and decreases down the group.
Across the table: Across the table: electron shells remain the same, but number of protons is increasing = protons have a greater pull on valence electrons, so are harder to remove.
Down the group: electron shells are increasing so there is electron shielding from the pull of the protons make the valence electron easier to remove.

Electron configuration of ions

When transition metals form ions they will lose electrons from the 4s suborbital first even though this orbital is filled first.

Example:
Chromium atoms and ions
[Cr] 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d5
[Cr2+] 1s2,2s2,2p6,3s2,3p6,3d5

electron configuration

Titration curves Part 2

Strong acid and weak base

Key points:

Initial pH: Ph = -log [H3O+]

No buffer region as it is a strong acid

Equivalence point: [H3O+] = SQ ROOT (ka [weak acid at equivalence])

Titration curves

Weak acid and strong base:

Key points:
Initial pH: [H3O+] = SQ ROOT( Ka [weak acid])
1/2 equivalence: pH = pKa
Equivalence point: mostly conjugate base present so use pOH
[OH-] = SQ ROOT (Kb [base at equivalence])

Buffer solutions

Calculating the pH of a buffer solution:

This can be calculated using the expression for Ka and the Ka value for weak acid used to make the buffer solution.

The dilution of the acid and base need to be taken into account before plugging the values into the Ka expression.

Example:

Remember: pH = -log [H3O+]

pH, pKa, pKb

Acid base equilibrium expressions

Kw = Ka x Kb

Converting between [H3O+] and pH, Ka and pKa, Kb and pKb

-log[Ka]

H3O+, Ka or Kb ---------------> pH, pKa or pKb

inv-log[pKa]

pH, pKa or pKb ----------------> H3O+ , Ka or Kb

calculating pH of a weak base

What is the pH of a 0.1 molL-1 solution of NH3? Given Pka(NH4+) = 9.25

Ka = inv log - 9.25
Ka = 5.62x 10^-10

Kb = Kw/ Ka
Kb = 1 x 10^-14 / 5.62 x 10^-10
Kb = 1.77 x 10^-5

NH3 + H2O -----> NH4+ + OH-
[NH3] = 0.1
[NH4+] [OH-] = X

Kb = [NH4+][OH-] / [NH3]

Kb = [x][x] / [0.1]

[x] = square root of [1.77 x 10^-5] x [0.1]

[x] = [OH-] = 1.33 x 10^-3

pOH = -log[OH-]

pOH = -log [1.33 x 10^-3]

pOH = 2.876

pH = 14 - 2.876

pH = 11.12

calculating pH of a weak acid

Strong vs weak acid....... dilute vs concentrated acid