Atomic radius decreases across the period. This corresponds to the increase in the number of protons and electrons. The electrons added occupy the same principal quantum shell (energy level). This increases the overall nuclear charge so now the protons are pulling on the outer shell electrons making the element more tightly bound.
Ionic radius increases across the period. Because the ratio of protons to electrons becomes greater, cations are smaller than their neutral atoms. Because of electron-electron repulsion and less effective shielding, anions are always larger than their neutral atoms. When elements are isoelectric (having the same number and configuration of electrons) the rule that the higher the nuclear charge the smaller the ion applies. The smallest will have the highest positive charge and will get larger until the highest negative charge.
Example:
Mg 2+ < Na+ < F- < O 2-
Smallest largest
Mg 2+ < Rb+ < Se 2- < Te 2-
Smallest largest
In this case Te is farther down the periodic table than Se so it will be larger. Mg has a higher charge than Rb so it will be smaller.
Melting point increases going across and maxes at silicon because the metallic bonds have a sea of electrons to overcome. Silicon is a giant covalent compound with a large number of contact points and van der waals forces to overcome. The nonmetals are simple covalent compounds which have weak van der waals forces to overcome.
Conductivity increases from Na to Mg to Al because they are giant metallic structure and the number of delocalized electrons contributing to the sea increases. Conductivity decreases and melting point increases at Silicon because it is a giant covalent structure. S, Cl & Ar are nonmetallic and have low BP due weak VDW forces
Electronegativity increases going across the period due to the increased number of protons attracting electrons.
Ionization energy increases across the period. Exceptions are group 3 < 2 due to the ease to remove the electron from the lone p orbital. 6 < 5 due to the ease to remove the electron from the doubly occupied px orbital.
Candidates should, for the third period (sodium to argon), be able to:
(a) *describe qualitatively (and indicate the periodicity in) the
variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)
(b) explain qualitatively the variation in atomic radius and ionic radius
(c) interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements
(d) explain the variation in first ionisation energy
(e) describe the reactions, if any, of the elements with oxygen (to give Na2O, MgO, Al 2O3, P4O10, SO2, SO3), chlorine (to give NaCl , MgCl 2, Al 2Cl 6, SiCl 4, PCl 5) and water (Na and Mg only)
(f) state and explain the variation in oxidation number of the oxides and chlorides in terms of their valance shell electrons
(g) describe the reactions of the oxides with water
[treatment of peroxides and superoxides is not required]
(h) describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids
(i) describe and explain the reactions of the chlorides with water
(j) interpret the variations and trends in (f), (g), (h), and (i) in terms of bonding and electronegativity
(k) suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties
In addition, candidates should be able to:
(l) predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
(m) deduce the nature, possible position in the Periodic Table, and identity of unknown elements from given information about physical and chemical properties
Ionic radius increases across the period. Because the ratio of protons to electrons becomes greater, cations are smaller than their neutral atoms. Because of electron-electron repulsion and less effective shielding, anions are always larger than their neutral atoms. When elements are isoelectric (having the same number and configuration of electrons) the rule that the higher the nuclear charge the smaller the ion applies. The smallest will have the highest positive charge and will get larger until the highest negative charge.
Example:
Mg 2+ < Na+ < F- < O 2-
Smallest largest
Mg 2+ < Rb+ < Se 2- < Te 2-
Smallest largest
In this case Te is farther down the periodic table than Se so it will be larger. Mg has a higher charge than Rb so it will be smaller.
Melting point increases going across and maxes at silicon because the metallic bonds have a sea of electrons to overcome. Silicon is a giant covalent compound with a large number of contact points and van der waals forces to overcome. The nonmetals are simple covalent compounds which have weak van der waals forces to overcome.
Conductivity increases from Na to Mg to Al because they are giant metallic structure and the number of delocalized electrons contributing to the sea increases. Conductivity decreases and melting point increases at Silicon because it is a giant covalent structure. S, Cl & Ar are nonmetallic and have low BP due weak VDW forces
Electronegativity increases going across the period due to the increased number of protons attracting electrons.
Ionization energy increases across the period. Exceptions are group 3 < 2 due to the ease to remove the electron from the lone p orbital. 6 < 5 due to the ease to remove the electron from the doubly occupied px orbital.
Candidates should, for the third period (sodium to argon), be able to:
(a) *describe qualitatively (and indicate the periodicity in) the
variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)
(b) explain qualitatively the variation in atomic radius and ionic radius
(c) interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements
(d) explain the variation in first ionisation energy
(e) describe the reactions, if any, of the elements with oxygen (to give Na2O, MgO, Al 2O3, P4O10, SO2, SO3), chlorine (to give NaCl , MgCl 2, Al 2Cl 6, SiCl 4, PCl 5) and water (Na and Mg only)
(f) state and explain the variation in oxidation number of the oxides and chlorides in terms of their valance shell electrons
(g) describe the reactions of the oxides with water
[treatment of peroxides and superoxides is not required]
(h) describe and explain the acid/base behaviour of oxides and hydroxides including, where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids
(i) describe and explain the reactions of the chlorides with water
(j) interpret the variations and trends in (f), (g), (h), and (i) in terms of bonding and electronegativity
(k) suggest the types of chemical bonding present in chlorides and oxides from observations of their chemical and physical properties
In addition, candidates should be able to:
(l) predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
(m) deduce the nature, possible position in the Periodic Table, and identity of unknown elements from given information about physical and chemical properties