The definition of electronegativity and how it is measured.
To apply our understanding of electrostatic principles to the periodic trends in electronegativity.
To predict the electronegativity of elements compared to each other.
How metallic and acid-base properties change across a period.
As seen in Periodic trends: Atomic radius, chemists have found, through experimenting, some principles of electrostatic forces forces that exist because charged particles attract or repel each other. Some principles are:
#1: Oppositely charged particles attract each other, while particles of like charge repel each other.
#2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force between a 2+ ion and a 2- ion is stronger than the attractive force between a 1+ and a 1- ion).
#3: Attractive forces between oppositely charge particles decrease with distance.
#4: Repulsive forces between like-charged particles decrease with distance.
Electronegativity is the ability of an atom (specifically the nucleus) to attract bonding electrons to its outer electron shell. It is measured using the Pauling scale fluorine is highest at 4.0 on the scale, the most electronegative element, whilst francium is the lowest at 0.7 and is the least electronegative element.
Electrostatic theory explains the trend in electronegativity in the Periodic Table, in the Periodic Table both across a period and down a group:
As you go across the elements in a period, each element has an increased effective nuclear charge, or Zeff attracting its outer shell electrons. Effective nuclear charge is the positive charge that the outer shell electrons effectively feel from the nucleus. To find this, subtract the number of inner electrons from the number of protons in the atom.
In lithium for example, Zeff = 3 - 2 = +1.
This equation means that two of lithiums 3 protons are cancelled out by the two inner shell electrons and only one is available for the one outer shell electron. Using electrostatic theory: Li is a Zeff 1+ nucleus attracting 1- of electron charge.
Moving across to beryllium, Zeff = 4 - 2 = 2+.
Using electrostatic theory: Be is a 2+ Zeff nucleus attracting 2 outer shell electrons or a 2- charge..
Increasing Zeff means the nucleus attracts and holds its outer shell electrons with increasingly greater force. It also better attracts other electrons to complete the outer shell.
Therefore, as you go across the period, it is easier for atoms to attract bonding electrons into their outer shell. This means electronegativity is higher.
As you go down a group of the periodic table, each element has an extra inner electron shell as well as increased nuclear charge. However, the increased nuclear charge is cancelled out by the extra inner shell, so Zeff is unchanged.
The extra inner shell of electrons also causes shielding of the nucleus. This is where the outer electrons are repelled away from the nucleus because of the like-charge inner shell electrons. In other words, positive attracts negative but not if there is a big wave of negative charge between them!
This means that going down the group, nuclei lose their ability to attract and hold onto bonding electrons. As a result, electronegativity decreases as you go down a group in the Periodic Table.
Remember: Noble gases are not given an electronegativity value because their atoms generally do not form bonds and since they already have full outer shells, they do not attract electrons to complete full outer shells!
The trends in electronegativity mean fluorine is the most electronegative element. The effect of electron shielding down a group is more influential than the effect of increased nuclear charge across a period, so oxygen is the second most electronegative element (around 3.5 on the Pauling scale), followed by chlorine (around 3.0).
The difference in electronegativity of atoms affects how different atoms bond with one another and can lead to substances of varying properties. This is a very important part of chemistry which the next chapter will look at – bonding between atoms and the properties of compounds they make as a result!>.
Electronegativity affects the metallic character of an element (metals have very low electronegativity) and as such, metallic character across a period changes. From left to right across a period, metallic character decreases. For example, in period 3:
Sodium, magnesium and aluminium are metals;
Silicon is a metalloid (it displays properties of both metals and nonmetals);
Phosphorus, sulfur, chlorine and argon are all nonmetals.
There is a sliding scale of metallic character across the period which affects several properties. One of these is the properties of oxide compounds. Using period 3 as an example:
Sodium oxide (Na2O), magnesium oxide (MgO) and aluminium oxide (Al2O3) are metal oxides with an ionic structure. As such, they usually have high melting points. When dissolved in water, metal oxides form basic solutions with a pH of over 7. The respective equations for sodium oxide and magnesium oxide reacting with water are:
Na2O + H2O → 2NaOH
MgO + H2O → Mg(OH)2
Sodium hydroxide is a strong base, which is why sodium oxide in water creates a basic solution. As you progress left to right, the metal oxides become less basic in nature.
Aluminium oxide is amphoteric; it acts a base reacting with acids and acts as an acid when reacting with bases, producing an aluminium salt both times. It also does not react with water in the way magnesium and sodium oxide does. Its reaction with acid is:
Al2O3 + 6HCl → 2AlCl3 + 3H2O
As silicon is a metalloid, silicon dioxide has very different properties to metal oxides like Na2O or MgO, in fact it is very weakly acidic. SiO2 has a giant covalent structure and does not react with water as the structure cannot be broken down by the interactions with H2O molecules. As it is weakly acidic, it will react with some metal oxides to produce silicon salts.
Phosphorus, sulfur and chlorine oxides are simple covalent molecules; comparing to metal oxides they have much lower melting and boiling points. When dissolved in water the nonmetal oxides form acidic solutions with a pH below 7. There are multiple oxides of phosphorus, sodium and chlorine but they all react to produce an acidic species with water.
The equation for phosphorus (V) oxide and water is:
P4O10 + 6H2O → 4H3PO4
Sulfur trioxide reacts to produce sulfuric acid, a reaction which is run on an industrial scale in the contact process:
SO3 + H2O → H2SO4
In the same manner, chlorine (VII) oxide reacts with water to produce perchloric acid which is a very strong acid.
Cl2O7 + H2O → 2HClO4
Chlorine also has the chlorine (I) oxide form, Cl2O which reacts with water to form a weaker acid, HOCl.
This trend of acidic nonmetal oxides is not only in period 3. Oxides of nitrogen like nitrogen dioxide, NO2, react with water to produce nitric acid, HNO3 and the weak acid nitrous acid, HNO2:
2 NO2 + H2O → HNO3 + HNO2
The nitrous acid breaks down over time into more nitric acid and NO gas. Like those in period 3, nitric acid is a strong acid and will produce very acidic solutions.
From all of this, you can determine that going from left to right across a period, oxide compounds go from being basic to acidic in nature.
The periodic trends of electronegativity
Periodic trends: Electronegativity
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