Chapter 8 - Periodicity Flashcards
Nickel is a metal with a high melting point.
State the block in the Periodic Table that contains nickel. (1)
D block
Give the full electron configuration of the Ni2+ ion (1)
1s2 2s2 2p6 3s2 3p6 3d8 (4s0)
Complete the electronic configuration of aluminium.
1s2 …………………………………………………………………………………………. (1)
1s2 2s2 2p6 3s2 3p1
State the block in the Periodic Table to which aluminium belongs. (1)
P block
The Ne atom and the Mg2+ ion have the same number of electrons.
Give two reasons why the first ionisation energy of neon is lower than the third ionisation energy of magnesium. (2)
- Mg+2 has more protons than Ne
- Mg+2 electrons closer to nucleus
By reference to all the atoms involved explain, in terms of electrons, how Na2S is formed from its atoms
- electron transfer between Na to S
- 1 electron from each Na atoms
Explain the meaning of the term periodicity as applied to the properties of rows of elements in the Periodic Table.
Describe and explain the trends in atomic radius, in electronegativity and in conductivity for the elements sodium to argon. (13)
Periodicity= trend in the properties of element across a period,
Elements in the same group have similar properties,
Atomic radius= decreases across the row, nuclear charge increases,
More attraction for electrons,
Electronegativity= increases across the row,
Nuclear charge increases,
Atomic radius decreases,
More attraction for shared electrons
Conductivity= decreases across the row,
Na to Al are metals,
Electrons are free to move in metals
Write an equation to illustrate the process occurring when the second ionisation energy of magnesium is measured. (1)
Mg+(g) = Mg^+2 (g) + 2e-
State and explain the trend in the melting points of the Period 3 metals Na, Mg and Al (3)
- increases,
- strong metallic bond,
- strong attraction between positive ions and delocalized electrons
Explain the meaning of the term first ionisation energy (1)
Energy required to remove one mole of 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions
State and explain the general trend in first ionisation energies for the elements Na to P (3)
- increases,
- nuclear charge increases,
- similar shielding,
- atomic radius decreases
State which one of the elements from Na to P has the highest melting point and explain your answer.
- silicon,
- giant covalent,
- strong covalent bonds need to be broken
State which one of the elements from Na to P deviates from this general trend and explain why this occurs (3)
- aluminum,
- electron in 3p orbital,
- less energy needed to lose electron
Explain why the melting point of sulphur, S8, is higher than that of phosphorus, P4 (1)
-van der waals forces
State and explain the trend in melting point of the Group II elements Ca–Ba. (3)
- decreases,
- weak metallic bond,
- increase in size of an atom
State the trend in atomic radius from phosphorus to chlorine and explain the trend (3)
- decreases,
- nuclear charge increases,
- similar shielding
In terms of structure and bonding, explain why sulfur has a higher melting point than phosphorus (3)
- sulfur molecules larger than phosphorus,
- van der waals forces between molecules are stronger,
- more energy needed to loosen forces between molecules
In terms of atomic structure, explain why the van der Waals’ forces in liquid argon are very weak (2)
- argon particles are single atoms with electrons closer to nucleus,
- cannot easily be polarized
Explain why nickel is ductile (can be stretched into wires). (1)
Layers can slide over each other
Explain, in terms of its structure and bonding, why nickel has a high melting point. (2)
Bonding= strong metallic bonds, Structure= strong attraction between positive ions and delocalized electrons
Describe the bonding in metals. (2)
- lattice of metal surrounded by delocalized electrons,
- strong attraction between positive ions and delocalized electrons
Explain why the melting point of magnesium is higher than that of sodium. (3)
- greater nuclear charge,
- stronger attraction between positive ions and delocalized electrons,
- Mg has more delocalized electrons,
- Na is a smaller atom
Explain why the first ionisation energy of sulphur is lower than would be predicted from the general trend (2)
- spin pair repulsion,
- repulsion between electrons in the 3p sub shell
Explain how the gaseous atoms of rubidium are ionised in a mass spectrometer (2)
- electron gun,
- knocks out electrons
State one reason why the first ionisation energy of rubidium is lower than the first ionization energy of sodium (1)
- Rb is a bigger atom,
- electrons are further away from the nucleus,
- more shielding in rubidium
By reference to the relevant part of the mass spectrometer, explain how the abundance of an isotope in a sample of rubidium is determined. (2)
- detector,
- current electrical signal related to abundance
