Unit 3
tructure 3.1: The periodic table: Classification of elements
S
.1.1:The periodic table: Classification of elements
3
- The periodic table consists of periods, groups and blocks.
- Identify the positions of metals, metalloids and non-metals in the periodic table.
3.1.2: Groups and periods of the periodic table
- The period number shows theouter energy levelthatis occupied by electrons.
- Elements in a group have acommon number of valenceelectrons.
- Deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa.
Note that
● Groups are numbered from 1 to 18.
● Only the following classifications should be known:
○ alkali metals
○ Transition elements (≠ Zn,Cd,Hg, and Cn)
○ Halogens
○ Noble gases
3.1.3: Periodicity
- Periodicity refers to trends in properties of elements across a period and down a group.
- Explain the periodicity of atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.
- Atomic radius
- Thedistance from the center of the nucleus to theoutermost electron shellis not measurable so itis half the
distance between two neighboring nuclei.
- Thisincreases down the groupbecause there are moreoccupied electron shells and therefore the atom is larger.
- Thisdecreases across the periodbecause there aremore protons in the nucleus from left to right in the same period
(a greater effective nuclear charge) This means the nucleus attracts the outer shell electrons more and so they
are pulled closer in.
, - Ionic radius
- This is measured in the same way as the atomic radius, using ions.
- Positive ions are smaller than their atoms because they have lost an electron shell and there is now a greater
effective nuclear charge pulling on the remaining electrons.
- Negative ions are larger than their atoms because they have gained extra electrons; there is now a smaller effective
nuclear charge pulling on the greater number of electrons.
- Thisincreases down the groupbecause there are moreoccupied electron shells and therefore the ion is larger
- From group I to IV this decreases for positive ions.All ions have the same electron arrangement but the effective
nuclear charge is greater due to more protons in the nucleus.
- From group IV to VII for negative ions this decreases.All ions have the same number of electrons but the effective
nuclear charge is greater due to more protons in the nucleus.
- Ionization energy
- X(g)
→X+ (g) + e-
- This is the amount of energy needed to remove one mole of electrons from one mole of gaseous atoms.
- Thisdecreases down the groupbecause the outermostelectrons are in a shell further out. This means there is a
lower attraction between valence electrons and the nucleus.
- Thisincreases across the periodbecause there aremore protons in the nucleus from left to right in the same period,
for the same number of electron shells. (the nucleus attracts the outer shell electrons more strongly and so they
are harder to remove.)
- Electron affinity
- X(g) + e-
→ X
-
(g)
- This is the energy change when one mole of electrons is gained by one mole of gaseous atoms. (Not given for noble
gases)
- This becomes(generally) less negative down the groupbecause the electrons are attracted into an energy level
further from the nucleus and this attraction is therefore not as strong.
- Thisgenerally becomes more negative across the period,as the effective nuclear charge increases and hence the
attraction for electrons into the valence energy level (but trend is a bit more complex)
- Electronegativity
- This is the ability of an atom to attract a pair of electrons shared in a covalent bond. It is a relative scale.
- Thisdecreases down the groupbecause the bondingpairs of electrons are further from the nucleus and there is
reduced attraction.
- Thisincreases across the periodbecause there aremore protons in the nucleus from left to right in the same period
(a greater effective nuclear charge.) This means the nucleus attracts pairs of electrons more strongly.
- Diagonal trend up to fluorine
3.1.4: Trends in properties
- Trends in properties of elementsdown a groupincludetheincreasing metallic character of group 1elementsanddecreasing
non-metallic character of group 17elements.
- Metallic character is the level of reactivity of a metal. It depends on the ease with which the metal loses an electron.
- Non-metallic character is the level of reactivity of a non-metal. It depends on the ease with which the non-metal gains an electron.
- Describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions.
tructure 3.1: The periodic table: Classification of elements
S
.1.1:The periodic table: Classification of elements
3
- The periodic table consists of periods, groups and blocks.
- Identify the positions of metals, metalloids and non-metals in the periodic table.
3.1.2: Groups and periods of the periodic table
- The period number shows theouter energy levelthatis occupied by electrons.
- Elements in a group have acommon number of valenceelectrons.
- Deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa.
Note that
● Groups are numbered from 1 to 18.
● Only the following classifications should be known:
○ alkali metals
○ Transition elements (≠ Zn,Cd,Hg, and Cn)
○ Halogens
○ Noble gases
3.1.3: Periodicity
- Periodicity refers to trends in properties of elements across a period and down a group.
- Explain the periodicity of atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.
- Atomic radius
- Thedistance from the center of the nucleus to theoutermost electron shellis not measurable so itis half the
distance between two neighboring nuclei.
- Thisincreases down the groupbecause there are moreoccupied electron shells and therefore the atom is larger.
- Thisdecreases across the periodbecause there aremore protons in the nucleus from left to right in the same period
(a greater effective nuclear charge) This means the nucleus attracts the outer shell electrons more and so they
are pulled closer in.
, - Ionic radius
- This is measured in the same way as the atomic radius, using ions.
- Positive ions are smaller than their atoms because they have lost an electron shell and there is now a greater
effective nuclear charge pulling on the remaining electrons.
- Negative ions are larger than their atoms because they have gained extra electrons; there is now a smaller effective
nuclear charge pulling on the greater number of electrons.
- Thisincreases down the groupbecause there are moreoccupied electron shells and therefore the ion is larger
- From group I to IV this decreases for positive ions.All ions have the same electron arrangement but the effective
nuclear charge is greater due to more protons in the nucleus.
- From group IV to VII for negative ions this decreases.All ions have the same number of electrons but the effective
nuclear charge is greater due to more protons in the nucleus.
- Ionization energy
- X(g)
→X+ (g) + e-
- This is the amount of energy needed to remove one mole of electrons from one mole of gaseous atoms.
- Thisdecreases down the groupbecause the outermostelectrons are in a shell further out. This means there is a
lower attraction between valence electrons and the nucleus.
- Thisincreases across the periodbecause there aremore protons in the nucleus from left to right in the same period,
for the same number of electron shells. (the nucleus attracts the outer shell electrons more strongly and so they
are harder to remove.)
- Electron affinity
- X(g) + e-
→ X
-
(g)
- This is the energy change when one mole of electrons is gained by one mole of gaseous atoms. (Not given for noble
gases)
- This becomes(generally) less negative down the groupbecause the electrons are attracted into an energy level
further from the nucleus and this attraction is therefore not as strong.
- Thisgenerally becomes more negative across the period,as the effective nuclear charge increases and hence the
attraction for electrons into the valence energy level (but trend is a bit more complex)
- Electronegativity
- This is the ability of an atom to attract a pair of electrons shared in a covalent bond. It is a relative scale.
- Thisdecreases down the groupbecause the bondingpairs of electrons are further from the nucleus and there is
reduced attraction.
- Thisincreases across the periodbecause there aremore protons in the nucleus from left to right in the same period
(a greater effective nuclear charge.) This means the nucleus attracts pairs of electrons more strongly.
- Diagonal trend up to fluorine
3.1.4: Trends in properties
- Trends in properties of elementsdown a groupincludetheincreasing metallic character of group 1elementsanddecreasing
non-metallic character of group 17elements.
- Metallic character is the level of reactivity of a metal. It depends on the ease with which the metal loses an electron.
- Non-metallic character is the level of reactivity of a non-metal. It depends on the ease with which the non-metal gains an electron.
- Describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions.
