T03/024 : ATOMIC PROPERTIES & CHEMICAL BONDING

Question 1

Draw resonance structures of carbonate and determine the most stable one

Question 2

Define resonance

Question 3

Give 4 characterisitics of resonance structures

Question 4

What is an isoelectronic molecule?

Question 5

What are the defining qualities of isoelectronic molecules?

Answer 5

• same number of valence electrons
• same geometry

Question 6

Define Lewis structures

Question 7

Identify the steps in writing Lewis structures

Question 8

What are the guidelines for Lewis structures

Question 9

Draw Lewis structure of PCl5

Question 10

Draw Lewis structure of SF6

Question 11

Define Formal Charge

Question 15

Define a chemical bond

Answer 15

chemical bond is the force of attraction between 2 or more atoms, ions or molecules that enable formation of compounds or molecules.

Question 16

Give 6 types of chemical bonds

Answer 16

1.Ionic or electrovalent bond
2.Covalent bond
3.Co-ordinate covalent bond
4.Metallic
5.Van der Waals
6.Hydrogen bonds

Question 17

What causes bonds to form?

Question 18

Define :
1. Valence
2. Valence electrons
3. Bonding Electrons
4. Non-bonding electrons
5. Electronic theory/Valence Theory/Octet theory of valence

Answer 18

• valence is the number of bonds formed by atom in a molecule
• valence electrons are the electrons in the outer energy level of an atom that can take part in chemical bonding
• Bonding electrons are electrons that are actively taking part in bond formation.
• non-bonding are still available for bond formation.
• Electronic theory or valence also called the Octet theory of Valence, states
  ‘Atoms interact by electron-transfer or electron-sharing, so as to achieve the
  stable outer shell of eight electrons.’

Question 19

State the octet theory of valence

Answer 19

Atoms interact by electron-transfer or electron-sharing so as to achieve the stable outer shell of 8 electrons.

Question 20

Differentiate the rule of 2 and the rule of 8

Answer 20

• the tendecy of atoms to have 8 e in their outer shell
• Since helium has two electrons in the outer shell, for hydrogen and lithium, having one and three (2, 1) electrons respectively, it is the Rule of two which will apply.

Question 21

Why do atoms form chemical bonds?

Answer 21

• this is to ensure the system achieves the lowest possible potential energy.
• Bond formation involves attraction forces overcoming repulsive forces between atoms and results in a net lowering of energy.

Question 22

Define an ionic bond

Answer 22

• electrostatic force of attraction between a cation and anion produced by electron transfer.

Question 23

State Coulomb’s Law and explain its relation to ionic bond

Answer 23

Coulomb’s Law and its relation to ionic bonding:

Coulomb’s Law states that:
F=k(Q1 x Q2)/r^2
k = (2.31 x 10^-19 J nm)

Where:
- F = force of attraction/repulsion between charges
- k = Coulomb’s constant
- q₁, q₂ = charges of the ions
- r = distance between the charges

Relation to Ionic Bonding:

1. Charge Effect:
   - Larger charges result in stronger ionic bonds
   - Example: Mg²⁺ forms stronger ionic bonds than Na⁺
   - This is why compounds like MgO have higher lattice energies than NaCl
2. Distance Effect:
   - Smaller distance between ions results in stronger bonds
   - Force of attraction increases as ions get closer
   - This is why smaller ions form stronger ionic bonds
3. Lattice Energy:
   - Directly related to Coulomb’s law
   - Higher charges and smaller ionic radii lead to greater lattice energy
   - The greater the lattice energy, the stronger the ionic bond

Think of it like magnets - the closer they are and the stronger their charges, the greater their attraction, just as ionic bonds become stronger with higher charges and smaller distances between ions.

Question 24

What can help you predict the stability of an ionic compund?

Answer 24

• charge of anion and cation

Question 25

What are the conditions necessary for formation of ionic bond?

Answer 25

**1) Number of valence electrons
*** Electropositive atom should have 1, 2 or 3
valence electrons while the electronegative
atom should have 5, 6 or 7 valence
electrons.
* Group 1, 2 and 3 satisfy condition for
electropositive atom, while group 5, 6 and
7 satisfy condition for electronegative
atom.

**2) Net lowering of energy
*** To form a stable ionic compound, there
must be a net lowering of energy.
* This means that energy must be released as
a result of electron transfer and final
formation of the ionic compound.

i) Low ionization energy of electropositive
element.
* The removal of electron from the
electropositive atom must requires energy
which should be low.
ii) High electron affinity for electronegative
atom.
* Addition of the electron to electronegative
must release high energy which is electron
affinity.
iii) High lattice energy.
* The electrostatic attraction between the
cation and anion in the solid compound
release energy which is electrical energy or
lattice energy.

N/B :- net energy should be negative

3.Electronegativity difference between metal and non-metal
- a large difference in electronegativity is needed for formation of ionic bond.
- Preferably a difference of 2 or more
is necessary for formation of ionic
bond between the two atoms.

Question 26

How can net lowering of energy be achieved during formation of ionic compunds?

Answer 26

1. The electropositive element should have a low ionization energy
2. The electronegative element should have high electron affinity
3. High lattice energy (large electrostatic attraction between cation and anion)

Question 27

What forces are at play in electron affinity? What leads to a positive electron affinity or negative electron affinity?

Question 28

What is lattice energy?Differentiate lattice energy formation and

Question 29

What are the factors governing the formaton of ionic bonds?

Question 30

Explain the factors that affect lattice energy

Question 31

look at charges

Between MgCl2 and CaO, which is the most stable ionic compound?What about between Al203 and CaO?Why?

Answer 31

The stability of ionic compounds depends on the lattice energy, which is influenced by the charge and size of the ions involved.

Comparison 1: MgCl₂ vs. CaO
- MgCl₂:
- Magnesium has a +2 charge, and chlorine has a -1 charge.
- The lattice energy depends on the attraction between Mg2+ and Cl- but Cl- relatively large ionic radius, reducing lattice energy.

• CaO:
  
  • Calcium has a +2 charge, and oxygen has a -2 charge.
  • The charges are higher compared to MgCl₂ O2- vs Cl-
  • Both ions in CaO are smaller than in MgCl₂, which increases the lattice energy due to stronger electrostatic attractions.

Conclusion: CaO is more stable than MgCl₂ due to its higher lattice energy, resulting from the higher charge and smaller ionic radii of Ca2+ and O2-

Comparison 2: Al₂O₃ vs. CaO
- Al₂O₃:
- Aluminum has a +3 charge, and oxygen has a -2 charge.
- The charges are higher, which significantly increases lattice energy.
- Both Al3+, O2- and Ca2+ and O2- further strengthening the electrostatic attraction.

Conclusion: Al₂O₃ is more stable than CaO because of the higher charge of Al3+, leading to much stronger lattice energy.

Summary:
1. CaO is more stable than MgCl₂.
2. Al₂O₃ is more stable than CaO.

This is due to the lattice energy being proportional to the product of the charges and inversely proportional to the sum of the ionic radii.

Question 32

Explain 4 charateristics of ionic compunds

Answer 32

• solid at room temperature
• High melting points
• Hard and brittle
• Soluble in water
• Good conuctors of electricity in..
• Do not exhibit isomerism
• ## Ionic reactions are fast

Question 33

What is the formula for the Born-lande equation?

Answer 33

                 N_A * M * z⁺ * z⁻ * e²
U = - --------------------------------------- * (1 - 1/n)
         4 * π * ε₀ * r₀

NA = Avogadro constant = 6.022x1023 atom-1
M = Madelung constant, relating to the geometry of the crystal;
Z+ = Charge on cation
Z- = charge on anion
e- = Electronic charge
ro = Sum of radii of cation and anion
n = Born exponent, typically a number between 5 and 12,
determined experimentally by measuring the compressibility of
the solid, or derived theoretically.
εo = 8.854 x10-12 J-1C2m-1

Question 34

What is the significance of Born-Lande equation?

Answer 34

Therefore, two key conclusions about factors that influence lattice
energy of ionic compounds:
1)From U α Z+Z-
The higher the charge on cation and anion, the greater the
magnitude of lattice energy
e.g. U (LiF) < U(CaF2) < U(MgS)

2) U α 1/ro
 The smaller the size of ions, the higher the lattice energy e.g.
U (NaF) > U(NaCl) > U(NaBr) > U(NaI)

Question 36

What is the Born-Haber cycle?

Question 37

State Hess’ law

Answer 37

The total amount of heat evolved or absorbed in a chemical reaction is constant whether the reaction is carried out in a single step or multiple steps.

Question 38

Calculate:
(I)lattice energy using Born-Haber cycle
(ii)Madelung constant using Borne- Lande equation

Question 39

ionic lattice, exothermix, endothermic, isolated gaseous ions

Define
(i)Lattice Dissociation Enthalpy
(ii)Lattice Formation Enthalpy

Answer 39

(i)Lattice Dissociation Enthalpy
(ii)Lattice Formation Enthalpy - the enthalpy change…

Question 40

How does charge density influence lattice enthalpy and melting points?

Question 41

Which substance in the the following pairs has the larger lattice
enthalpy?
a) NaCl or KCl b) NaF or NaCl
c) MgCl2 or NaCl d) MgO or MgCl2

Answer 41

(a) KCl
(b)NaF
(c)MgCl2
(d) MgO

Question 42

charge, lattice enthalpy

Magnesium oxide is used to line furnaces. Why?

Question 43

Explain trends of thermal stability and decomposition down Group 2

Question 44

Define and state whether they are exothermic and endothermic :
1.Standard Enthalpy change of formation
2.Standard Enthalpy Change of Atomisation
- enthalpy of sublimation
-
3.First Ionisation energy
4.Second ionisation energy
5.Electron affinity

(if you’re using 1 atom of a molecule divide the energy of atomisation by no. of atoms in that molecule)

Question 45

Why are lattice compunds formed at a lower energy state than the reactants

Question 46

Calculate the ionisation energies of Mg^2+
(Calculate the change in energy in formation of Mg^2+ and Al^3+)

Question 47

Differentiate hydration and lattice enthalpy

Answer 47

for ionic compund to dissolve, hydrattion > lattice

Question 48

Define enthalpy of solution

Question 49

Define enthalpy of hydration

Answer 49

• multiply by no of ions formed

Question 50

Define solvent-solvent intermolecular attractions

Question 51

Describe the relationship btn The enthalpy of solution, ΔHsolution, enthalpy of hydration,
ΔHhydration, lattice energy and solvent-solvent intermolecular
attractions, ΔHosolvent-solvent

Answer 51

ΔH°sol = ΔH°LE + (ΔH°hyd(anion) + ΔH°hyd(cation)) + ΔH°solvent-solvent

Question 52

Explain how these factors affect solubility :
1. Solvent
2. Lattice energy
3. Enthalpy of hydration
4. Enthalpy of solution
5.

Question 53

what is the Effect of the Size and charge of ions on LE and ΔHhydration:

Answer 53

• compounds with smaller size have higher hydration enthalpy …

Question 54

Explain the trends in Lattice energy and melting points in : MgCl2, CaCl2, Mg0 and NaF, NaCl, NaBr

Question 55

What is the lattice energy for CaCl2 given the following thermodynamic
data.
a) Ca2+(g) + 2Cl-(g) →CaCl2(s)
b) sublimation energy of calcium = +178 kJ/mol
c) first ionization of calcium = +590 kJ/mol
d) second ionization of calcium = +1,150 kJ/mol
e) bond dissociation of Cl2(g) = +244 kJ/mol
f) electron affinity of Cl- = -349 kJ/mol
g) enthalpy of reaction or formation = -326 kJ

Question 56

If the lattice energy increases with charge and the association of
oppositely charged ions is a stabilizing force, can NaF form, with Na+2 and
F-2?

Question 57

Use the thermodynamic data of CuF given below and
other information from inorganic chemistry to
determine the value of Madelung of CuF. Energy is in
kJ/mol and eV.
Cu(s)→Cu (g) = 339.
Cu(g)→Cu+
(g) + e− = 745.
F2(g)→2F(g) = 243.4.
F(g) + e−→F−
(g) = -349.
Cu+
(g)→Cu2+
(g) + e =1960.
Cu(s) +½ F2 (g) →CuF (s) = -140.
Give your final answer to 2 decimal places.

Question 58

Describe the valence bond theory

Question 59

What are the main features of the valence bond theory?

Question 60

What are the different bond types?