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1.
Formation of ionic bonds
The concept of ionic bond was proposed by W.
Kossel in Germany in 1916
.
Atoms gain and lose electrons to form positive and negative ions, and the positive and negative ions are combined by electrostatic attraction to form ionic bonds
.
The electronegativity difference between the elements that can form ionic bonds is relatively large (generally △X>1.
7), and only a few electrons can be transferred to achieve a stable structure of the rare gas electron configuration
.
E.
g
E-na - → na +
+ E cl - → cl -
Ionic crystals are stable and release more energy when forming ionic compounds
.
Na+1/2Cl 2 =NaCl △rHm Θ =-411.
2kJ·mol -1
2.
The nature of ionic bond
Ionic bonds have no directionality and saturation
.
Ions can attract ions with opposite charges in any direction, so the ionic bond has no directionality; if space conditions permit, each ion attracts as many ions with opposite charges as possible to reduce energy
.
The research results show that the ionic compound is not purely electrostatic, and there is also some overlap of atomic orbitals (covalent bond components)
.
People use the percentage of ionicity to express the ionic size of the ionic bond.
If the percentage of ionicity in a compound exceeds 50%, it is considered that an ionic compound is formed
Table 6-1 shows the relationship between the percentage of ionicity and the difference in electronegativity when the AB type compound is combined with a single bond
.
Table 6-1 The relationship between the ionicity percentage of a single bond and the difference in electronegativity
Ionic bonds are generally formed between the atoms of active metals and active non-metallic elements with large difference in electronegativity
.
It is generally believed that the difference in electronegativity between the two elements △X>1.
7 will form an ionic compound, otherwise it will form a covalent compound
.
3.
Ionic bond strength
1) Measurement of ionic bond strength
People often use bond energy and lattice energy to measure the strength of ionic bonds
.
Refers 1mol bond energy is gaseous molecular gaseous dissociation atom, a 1mol disconnect the energy required for a chemical bond, represented by E
.
For diatomic molecules, the bond energy (E) is equal to the dissociation energy (D), such as Cl 2 molecules; for polyatomic molecules, the bond energy is equal to the average dissociation energy, such as NH 3 molecules
.
Lattice energy refers 1mol ionic crystal (crystal form ionic bonds) dissociation energy of positive and negative ions needed for gaseous, indicated by the U-
.
The monolithic crystals of ionic compounds can be regarded as giant molecules.
The binding force is not just the combination between positive and negative ions.
The lattice energy can not be directly measured, but it can be obtained indirectly from the calculation of thermodynamic data
.
Born and Haber designed a thermodynamic cycle to obtain the lattice energy from thermodynamic data, which is called the Born-Haber cycle
.
△H 1 is the heat of atomization of Na, △H 1 =108kJ·mol -1 ;
△H 2 is 1/2 of the dissociation energy of Cl 2 , △H 2 =121kJ·mol -l ;
△H 3 is the first ionization energy of Na, △H 3 =496kJ·mol -1 ;
△H 4 is the opposite of the electron affinity of Cl, △H 4 =-349kJ·mol -1 ;
△H 5 =-U, U is the lattice energy of NaCl;
△H 6 is the molar heat of formation of NaCl, △H 6 =-411kJ·mol -1 ;
According to Hess Law
△H 6 =△H 1 +△H 2 +△H 3 +△H 4 +△H 5
Then the lattice energy of NaCl is
U=-△H 5 =△H 1 +△H 2 +△H 3 +△H 4 -△H 6 =787(kJ·mol -l )
2) Factors affecting the strength of ionic bonds
The greater the electrostatic attraction between the positive and negative ions, the stronger the ionic bond, the higher the boiling point and hardness of the ionic crystal, and the greater the lattice energy or bond energy
.
If the distance between two ions is d, and the charges of the positive and negative ions are q + and q -respectively , the potential energy V between the positive and negative ions is
In the formula, ε 0 is the relative permittivity
.
The charge of the ion affects the strength of the ionic bond
.
The higher the charge of the ion, the greater the attraction between the positive and negative ions, and the stronger the ionic bond
.
For example, melting point: NaCl is 801°C, MgO is 2825°C; lattice energy: NaCl is 787 kJ·mol -1 , MgO is 3916 kJ·mol -1
.
The radius of the ion affects the strength of the ionic bond
.
The smaller the radius of the ion, the smaller the distance d between the ions, the greater the electrostatic attraction between the positive and negative ions, and the stronger the ionic bond
.
For example, NaI a melting point of 660 ℃, much lower than the NaCl, because the I - radius ratio C - larger; Meanwhile, NaI lattice energy (686kJ · mol -1 ) is much smaller than NaCI
.
The electronic configuration of ions affects the strength of ionic bonds, and this effect is more complicated
.
The more the number of outer electrons of an ion, the higher the effective nuclear charge, the stronger the ionic bond; at the same time, the more the number of outer electrons of an ion, the more the ion’s polarization and deformability, and the covalent components between positive and negative ions increase.
, The ionic bond strength decreases
.
Experimental data shows that the latter factor is the main factor, that is, the more electrons in the outer layer of an ion, the lower the strength of the ionic bond
.
For example, the melting point of the compound: CaCl 2 (775°C)>MnCl 2 (650°C), CaO (2613°C)>MnO (1842°C), CaSO 4 (1460°C)>MnSO 4 (700°C)
.
Related links: Ion electron configuration and radius