(All you need to know about Ionic Bonding)
Guiding question: What determines the ionic nature and properties of compounds?
Table of Contents
Metal loses e– → cations (positive ions)
Non-metal gains e– → anions (negative ions)
Effective nuclear charge
↳ The presence of inner electrons/shielding electrons and the longer distance between the nucleus and the valence electron level reduces the attraction of the nucleus on the outer electrons. Hence, the outer electrons experience an effective nuclear charge which is less than the full nuclear charge.
↳ Formula:
Zeff = total number of protons – inner electrons
↳ Left to right in a period, one proton is added to the outer energy level. The effective nuclear charge increases as there is no change in the number of inner electrons. This effect also reduces the atomic size.
↳ In a group, the increase in the nuclear charge is offset (cancelled) by the increase in the number of inner electrons it remains approximately the same.
Metal atoms have low ionisation energies as ionisation energies decrease down a group, hence metals on the bottom left have greater tendencies to lose electrons and form ionic compounds.
Non-metals form negative ions as they have high effective nuclear charges so they attract transferred electrons more strongly.
You know what? Let’s look at an example and make this easier for the both of us to understand better.
Take the compound NaCl (Sodium Chloride)
↳ The outer 3s electron from a sodium atom moves to the chlorine atom, the resulting electrostatic attraction between Na+ and Cl– pulls the ions together and forms an ionic bond
mOviNg oN
Extra info y’all lols
- Si4+ does not exist in chemical reactions as the first ionisation energy of silicon is too high to make the electron loss.
- Negative ions are more attractive with increased effective nuclear charge.
- The addition of electrons becomes more difficult with increasing negative charge because of e– – e– repulsion.
- Si4- is not feasible for a reaction.
- Noble gases have very high ionisation energies and complete energy levels which makes them stable and unreactive. Hence, they do not gain or lose electrons.
- The effective nuclear charge of noble gas is 0.
- The number of outer/valence electrons helps to predict the position of an element in the periodic table.
HL
- Transition elements form ions of different charges and have variable oxidation states.
- The loss of electrons from 3d orbital is easier for the electron loss and hence, the transition elements exist in variable oxidation states. The ionisation energies of transition elements exist from +1 to +7.
- Ions with a charge greater than 3+ generally show covalent character as they have large charge density, they can polarise negative ions.
- The high charge density of the transition elements and their ions pull the weakly held outer electrons. This is called polarisation of negative ions and it increases the covalent ions and it increases the covalent character of the transition metal with oxidation state > +1.
HELO SO LIKE JUST A LITTLE SOMETHING TO ADD HERE:
If you see 1+ anywhere then that means it’s the charge.
If you see +1 anywhere then that means it’s the oxidation state.
(I’ll elaborate on oxidation states in a later chapter)
SL
The attraction between ions increases with ionic charge. Ions of greater charge are formed if more than one electron is transferred.
The most reactive metal Caesium has the lowest ionisation energy and loses electrons easily.
Fluorine is the most reactive halogen with the smallest atomic radius and attracts the transferred electrons more strongly. The reactivity of halogens are sometimes used interchangeably with electronegativity, i.e. a halogen that is more electronegative is more reactive.
Naming of Compounds
Metal (charge if necessary) + acid radical/oxide
Include charge if variable charge
Polyatomic Ions
So like bad news guys. You gotta remember a lot of these ions and their charges and all that cause it’s not there in the data booklet 🙁
Here’s a list but I’m highlighting the ones that are needed for this chapter and in a few other chapters later on, but you would need to remember all of them at one point for your IB exams.
One way of remembering them effectively is to think of them as acid radicals, i.e. polyatomic ions are basically acids without the H+. Do note that there are exceptions: H2SO4, H3PO4, H3PO3, etc. which are variable.
Ionic Lattice
A lattice is a bonding structure that consists of very large numbers of ions and can grow indefinitely, their formula is basically the ratio of the ions present. It is called its empirical formula or formula units.
HL
Coordination number
It is used to express the number of ligands surrounding a given ion in the lattice.
Back to SL
Ionic Bonds are very strong due to electrostatic forces of attraction between opposite charges.
↳ Na+ + Cl– → NaCl
↳ ∆H = -ve (exothermic)
NaCl → Na+(g) + Cl–(g)
↳ hypothetically breaks to gaseous ions
Properties of an Ionic Lattice
- Lattice enthalpy is the amount of energy required to separate the ions. The value of enthalpy is positive.
The enthalpy change is calculated from the ionic model; it assumes that the crystals are made up of spherical ions interconnected by electrostatic forces.
Lattice enthalpy is greater for ions with large charge density as they have small radius, high charge and vice versa. Lattice enthalpy measures the strength of an ionic bond.
HL
– K is a constant that takes account of many ion interactions and also depends on the geometry of the lattice.
– n and m are the magnitudes of the charges.
– R is the radii of the ions. - Melting Point and Boiling Point (Volatility)
– Ionic compounds have high melting and boiling points as a large amount of energy is needed to separate the ions.
– The melting point generally increases when the ionic churches are greater
eg. NaF 1+ 1-
Na2O 1+ 2-
The melting point of Sodium Oxide is more than Sodium Fluoride as doubling the ionic charges increases the attraction between ions
eg. MgF2 2+ 1-
MgO 2+ 2-
MgO has higher melting point than MgF2
– The decrease in the melting point of the sodium halides is explained by the increase in ionic radius as attraction between the ions decreases
– Volatility is a tendency of a substance to vaporise or to change states of matter
– Ionic compounds have no volatility dues to high boiling points. Hence, they are non-volatile. - Solubility
– Ionic compounds are soluble in water but not in non-polar solvents.
– Water is a polar molecule and has small partial positive charges on the hydrogen atoms.
– Hydrogen atoms attract Cl– and small partial -ve charge on the oxygen atom attracts sodium ions, hence water molecules pull the ions from their lattice. Ions are now surrounded by water molecules and called as hydrated, and the solid dissolves.
– For solvents other than water, the ions are called solavated (solvation).
– For non-polar solvents there is no attraction between molecules and ions, they remain in their lattice and are insoluble.
– Exception → CaCO3
NaCl(s) → NaCl(aq)
Nacl(s) → Na+(aq) + Cl–(aq) - Electrical conductivity
– Ionic compounds do not conduct electricity in their solid state as ions are fixed in their lattice, but they can conduct electricity in their liquid or aqueous states as ions are mobile. - Brittle
– Ionic compounds are brittle. The crystals tend to shatter when a force is applied. The force displaces some ions in the lattice, which results in ions of the charge to rub alongside each other.
– Their brittle characteristics are generally due to the lattice structure as well as the nature of the bond
– The repulsive forces of the ions cause the lattice to split. - Period 3 chlorides are less ionic across the period.
NaCl – MgCl2 – AlCl3 – SiCl4 – PCl5 – SCl2 – Cl2
→ Melting point decreases →
Ionic → Covalent
– The transition of ionic to covalent character can be explained in terms of bonding. Sodium Chloride is ionic. MgCl2 is ionic. AlCl3 is considered to be covalent (EXCEPTION). The chlorides of the later elements are molecular covalent and have very low melting points.
→ molecular covalent: have fixed formulas, giant covalent don’t.
LAST TOPIC OF THIS CHAPTER LES GOOOOO
Ionic character and electronegativity
Electronegativity is the measure of the ability of an atom to attract electrons. That is quantified using the Pauling Scale.
↳ Fluorine has the most electronegative value of 4.
↳ Caesium has the most electropositive value.
↳ CsF is considered to be 100% ionic
↳
↳ A difference of 1.8 is considered predominantly ionic.
↳ E.D > 1.8 is ionic
↳ E.D < 1.8 is covalent
Electronegativity increases across a period and decreases down a group.
Past Paper Questions (SL & HL)
ans: C
explanation – The electrostatic forces of attraction are stronger and the lattice enthalpy is higher when the ions have a larger charge. The electrostatic attraction between oppositely charged ions is stronger and the charge density is higher for smaller ions. Statement one would not be true since the smaller the lattice enthalpy, the weaker the electrostatic attraction and the lower the melting point is.
ans:
a) Ionic bonding. Electrostatic attraction of (metal) cations and (non-metal) anions held together in a (crystal) lattice
explanation – An anion and a cation are produced when atoms with a significant difference in electronegativity, such as metals and non-metals, exchange electrons. This process is known as an ionic bond. An ionic connection is created when these ions are attracted to one another electrostatically. In this instance, the metal calcium loses electrons to the non-metal phosphorus.
b) Ca3P2
c) -1
explanation – Calcium has a charge of +2 in ion form. Therefore to form Ca(SCN)2, the charge on the thiocyanate must be -1
d) the larger the lattice enthalpy, the stronger the bond. We get this because the magnitude is the measure of the strength of the ionic bond.
explanation – Lattice enthalpy measures the strength of the forces between the ions in a crystal lattice. Therefore, a higher lattice enthalpy corresponds to a stronger bond. Lattice enthalpy is influenced by both the size of the ions and the ion charge. Smaller ions and higher charges result in stronger bonds and, therefore, higher lattice enthalpies.
e) O2- has a higher charge than F– and higher charges result in greater lattice enthalpy.
explanation – Stronger electrostatic interaction between the ions at higher charges leads to bigger lattice enthalpies. Coulomb’s Law, which states that the force between ions is inversely proportional to their distance apart and proportional to the product of their ion charges, can be used to explain this.
f) Calcium Oxide (CaO) will have a higher melting point because of it having larger lattice enthalpy / stronger ionic bond / greater force of attraction between ions.
explanation – The electrostatic attraction between the ions needs to be overcome in order to melt an ionic compound. Since the lattice enthalpy increases with bond strength, breaking a bond requires more energy.
Calcium oxide (CaO) will have a greater melting point than calcium fluoride (CaF2) due to its higher lattice enthalpy.
ans: B
explanation – In order to have a complete outer shell, magnesium has to lose both of its two electrons from its outermost shell. The outermost shell of fluorine has seven electrons; a complete outer shell would require one more electron. Therefore, to make a neutral ionic compound—where the charges of the positive and negative ions cancel out—two fluorine atoms are needed for every magnesium atom.
ans:
a)
i. LiCl has ionic bonding while HCl has polar covalent bonding. Electrostatic forces of attraction are stronger in ionic bonds than in polar covalent bonds. (polar covalent bonds have london-dispersion and dipole-dipole interactions)
ii. CaBr2 has ionic bonding while Bromine has nonpolar covalent bonding. Soluble/molten ionic compounds have free moving charged particles and can conduct electricity while covalent compounds do not have charged particles in their composition. So they cannot conduct electricity.
iii. HI has polar covalent bonds while I2 has nonpolar covalent bonding (found based on the electronegativity difference). The london-dispersion forces in I2 are much stronger than the dipole-dipole interactions of HI.
explanation to why LDF is stronger than dipole-dipole interactions in this question:
1. the molar masses – HI is 127.91 gmol−1 while I2 is 253.80 gmol−1. The LDF outweighs the dipole-dipole interactions.
2. Since the iodines in I2 have more electrons and a larger radius, the electrons in the outer valence shell are more free to move around resulting in greater instantaneous dipole moments meaning stronger dispersion forces.
b) Bromine < Potassium < Potassium Bromide
Br2 has pure covalent bonding; potassium has metallic bonding; potassium bromide has ionic bonding.Ionic bonding in potassium bromide results in the greatest intermolecular interactions.Even though potassium has metallic bonding, it only has one valence electron and covalent bonds have the weakest intermolecular forces.
WE’RE DONEEEEEEEEEE
Happy studying and ace those exams! Catch you later!
external resources used:
1. my chem notebook
2. Revision Village – 2.1 Ionic Bonding Question Bank
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