Periodicity Flashcards
Arrangement of the periodic table
elements placed in order of increasing atomic number
vertical columns called groups
horizontal rows called periods
alkali metals, halogens, noble gases, transition metals, lanthanoides and actnoides
non-metals
found on the upper right hand side of the p block
e.g. halogens are reactive non-metals
metallic elements
found on the left hand side of the table in s block, in the central d block and island of the f block
e.g alkali metals are a reactive group of metals in group 1 of the s block
the lanthanoides and actinoides are metals which make up the first and second row of f block
Metalloid elements
have the characteristics of both metals and non-metals
physical properties and appearance most resemble the metals, although chemically they have more in common with the non-metals
in the periodic table the metalloid elements silicon, germanium, arsenic, antimony, tellurium and polonium form a diagonal staircase betwene the metals and non-metals
Nuclear charge
is given by the atomic number and so increases by one between successive elements
Effective nuclear charge
the outer electrons that determine many of the physical and chemical properties of the atom do not experience the full attraction of this charge as they are shielded from the nucleus and repelled by inner electrons
hence the ‘effective charge’ experienced by the outer electrons is less than the full nuclear charge
effective nuclear charge = # protons - #electrons in inner shells
Trend in effective nuclear charge
increases across a period
remains approx +1 down the group as the increase in nuclear charge is offset by the increase in # of inner electrons
Atomic radius
measured as half the distance between neighbouring nuclei
(can also be considered as the distance from the nucleus to the outermost electrons)
Trend in atomic radius
increases down a group
as the number of occupied electron shells increases
decreases across a period
the attraction between the nucleus and outer electrons increases as the nuclear charge increases so there is a general decrease in atomic radii across the period
Trends in ionic radius
- positive ions are smaller than their parent atoms (loss of outer shell) increased attraction force between nucleus and electrons pulls the shells in
- negative ions are larger than parent atoms (adds electrons to the outer shell, increasing repulsion causing the electrons to move further apart, increasing radius of the outer shell)
- ionic radii decreases from groups 1 to 14 for the positive ions
- ionic radii decrease from groups 14 to 17 for the negative ions
- ionic radii increase down a group as the number of electron energy levels increases
First ionization energy
measure of the attraction between the nucleus and the outer electrons
provide evidence for the electron config of the atoms
Trend in ionization energy
- increase across a period. Increase in effective nuclear charge causes an increase in the attraction between the outer electrons and the nucleus and makes the electrons more difficult to remove
- decrease down a group. Although nuclear charge increases the effective nuclear charge is about the same, owing to shielding of the inner electrons and so the increased distance between the electron and the nucleus reduces the attraction between them
- small departures from these trends provide evidence for sub levels
- reverse of the trend in atomic radii
Electron affinity
is the energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions
first ea is exothermic and second ea is endothermic
Trend in electron affinity
the minimum values for electron affinities occur for the group 17 elements
incomplete outer energy level and high effective nuclear charge and so attract electrons the most
max for group 18 and 1 elements
lowest effective nuclear charge and so attract the extra electron the least
Electronegativity
measure of the ability of its atoms to attract electrons in a covalent bond
Trend in electronegativity
increases across a period owing to increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons
decreases down a group, the bonding electrons are furthest away from the nucleus and so there is reduced attraction
Melting points
depend on both the type of bonding and the structure
Trends in melting points
decrease down group 1 (metals held together by attractive forces between delocalised e and the positively charged ions, attraction decreases with distance)
increase down group 17 (molecular structures held together by LDF, increase with number of e)
generally rise across a period
Chemical properties
are determined by the electron configuration of its atoms
elements in the same group have similar chemical properties as they have the same number of valence electrons
Chem React. Group 1: the alkali metals
silvery metals - too reactive to be found in nature
react with water to produce hydrogen and the metal hydroxide
this reaction becomes more vigorous as the group is descended (most reactive is caesium)
Properties Group 17: halogens
exist as diatomic elements
physical:
- coloured
- show gradual change from gases to liquid to solids
chemical:
- very reactive non-metals
- reactivity decreases down the group
- form ionic compounds with metals and covalent compounds with other non-metals
Group 17 reaction with Group 1 metals
halogens react with group 1 elements to form ionic halides
halogen atom gains one e from group 1 element to form a halide ion X-
resulting ions both have the stable octet of the noble gases
the most vigorous reaction occurs between the elements which are the furthest apart (bottom group 1 and top group 17)
Group 17 displacement reactions
chem reactivity can be seen by placing them in direct competition for an extra electron
when Cl is bubbles through a solution of KBr the solution changes from colourless to orange owing to the production of Br2
2KBr + Cl2 -> 2KCl + Br2
2Br- + Cl2 -> 2Cl- + Br2
other reactions are:
2I- + Cl2 -> 2Cl- +I2
the colour changes from colourless to dark orange/brown owing to the formation of iodine
2I- + Br2 -> 2Br- + I2
The halides
halogens form insoluble salts with silver
adding a solution containing the halide to a solution containing silver ions produces a precipitate that is useful in identifying the halide ions
Ag+ + X- -> AgX
AgCl = white
AgBr = cream
AgI = yellow
Bonding of Period 3 oxides
oxides of elements Na to Al have giant ionic structures
oxides of P, S and Cl are molecular covalent
Si has a giant covalent structure
the ionic character of a compound depends on the difference in EN between its elements
oxides become more ionic down a group as the EN decreases
Basic oxides
oxides of metals are ionic and basic
sodium oxide and magnesium oxide dissolve in water to form alkaline solutions owing to the presence of OH- ions
Na2O + H2O -> 2NaOH
MgO + H2O -> Mg(OH)2
a basic oxide reacts with an acid to form a salt and water, the oxide combines with two H+ ions to form water
O2- + 2H+ -> H2O
Li2O + 2HCl -> 2LiCl + H2O
MgO + 2HCl -> MgCl2 + H2O
Acidic oxides
the non-metallic oxides react readily with water to produce acidic solutions
phosphorus (V) oxide reacts with water to produce phosphoric (V) acid
P4O10 + 6H2O -> 4H3PO4
phosphorous (III) oxide reacts with water to produce phosphoric (III) acid
P4O6 + 6H2O -> 4H3PO3
Amphoteric oxides
Aluminium oxide does not affect the pH when it is added to water as it is essentially insoluble
however, it shows both acid and base behaviour
e.g. behaves as a base as it reacts with H2SO4
Al2O3 + 6H+ -> 2Al3+ + 3H2O
Al2O3 + 3H2SO4 -> Al2(SO4)3 + 3H2O
and behaves like an acid when it reacts with alkalis such as NaOH
Al2O3 + 3H2O + 2OH- -> 2Al(OH)4-
D block elements - electron config
the similarity in the properties of the first d-block elements is illustrated by the relatively small range in atomic radii
correspondingly small increase in effective nuclear charge experiences by the 4s electrons
unusual config of chromium and copper are due to the stability of the filled and half filled 3d sub level
similarity in atomic radii explains ability of transition metals to form alloys (atoms of one d block metal can be replaced by atoms of another without too much disruption of the solid structure
Physical properties of d block elements
- high electrical and thermal conductivity
- high mp
- malleable
- high tensile strength
- ductile
- Fe, Co and Ni are ferromagnetic
these props. can be explained in terms of the strong metallic bonding found in the elements
Chemical properties of d block elements
- form compounds with more than one oxidation number
- form a variety of complex ions
- form coloured compounds
- act as catalysts
Transition metals
are elements whose atoms have an incomplete d sub shell, or which an give rise to cations with an incomplete d sub shell
zinc is not a transition metal (it is a d block element) but the d sub level is complete in both the atom and the ion
Explanation or variable oxidation number of transition elements
all the transition metals show both the +2 and +3 oxidation states
the M3+ ion is the stable state for the elements from scandium to chromium, but the M2+ state is more common for the later elements
the increased nuclear charge of the later elements makes it more difficult to remove a third electron
the maximum oxidation state of the elements increases in steps of +1 and reaches a maximum at manganese
these states correspond to the use of both the 4s and 3d electrons in bonding
thereafter the maximum oxidation state decreases in steps of -1
oxidation states above +3 generally show covalent character
Complex ions
transition metals in solution have a high charge density and attract water molecules which form coordinate bonds with the positive ions to form a complex ion
more generally a complex is formed when a central ion is surrounded by molecules or ions which possess a lone pair of e
these surrounding species (ligands), are attached via a coordinate bond. All ligands have at least one atom with a lone pair of electrons which is used to form a coordinate bond with the central metal ion
the number of coordinate bonds from the ligands to the central ion is called the coordination number
in aqueous solution, water molecules generally act as ligands
Polydentate ligands act as chelating agents
some species have more than one lone pair available to form a coordinate bond with the central transition ion
Transition metals and their ions are important catalysts
Fe in the haber process
N2 + 3H2 <=> NH3
Ni in the conversion of alkenes to alkanes
Pd and Pt in catalytic converters
MnO2 in the decomposition of hydrogen peroxide
V2O5 in the contact process
2SO2 + O2 <=> 2SO3
Sulfur trioxide is used in the production of sulfuric acid
Magnetic properties of the transition metals and their compounds
every spinning electron in an atom or molecule can behave as a tiny magnet
most substances have paired e that pair up and so are non-magnetic
some transition metals and their compounds are unusual in having some electrons that remain unpaired which when aligned lead to magnetic properties
Diamagnetism
is a property of all materials and produces a very weak opposition to an applied magentic field (paired e)
zinc
Paramagnetism
which only occurs with substances which have unpaired electrons, is stronger than diamagnetism. It produces magnetism proportional to the applied field and in the same direction.
property of single atoms or ions with unpaired spinning eletrons
transition metal complexes with unpaired electrons show paramagnetic properties as they are pulled into a magnetic field
increases with the number of unpaired e so generally increases from left to right across periodic table
Ferromagnetism
is the largest effect, producing magnetisations sometimes in orders of magnetism greater than the applied field
Fe, Ni and Co are ferromagnetic
the unpaired d e in large numbers of atoms line up with parallel spins in regions called domains
Coloured compounds
the d sub level splits into two sets of orbitals of different energy in a complex ion
complexes of d block elements are coloured as light absorbed when an electron is excited between the d orbitals
the colour absorbed is complementary to the colour observed
the colour of transition metal ions can be related to the presence of partially filled d orbitals (Sc3+ is colourless because the 3d sublevel is empty and Zn2+ is also colourless because the 3d sublevel is full)
what does colour of compound depend on
- nuclear charge and the identity of the central metal ion
- the charge density of the ligand
- the geometry of the complex ion (the electric field created by the ligand’s lone pair of e depends on the geometry of the complex ion)
- the number of d electrons present and hence the oxidation number of the central ion
Charge density of the ligand -> colour
the strength of the coordinate bond between the ligand and the central metal ion depends on the electrostatic attraction between the lone pair of e and the nuclear charge of the central ion. Ligands interact more effectively with the d orbitals of ions with a higher nuclear charge.
Spectrochemical series
arranges ligands according to the energy separation
the wavelength at which maximum absorbance occurs decreases with the charge density of the ligand
Geometry of complex -> colour
the splitting of energy of the d orbitals depends on the relative orientation of the ligand and the d orbitals
number of d e and oxidation state of the central metal ion -> colour
strength of the interaction between the ligand and the central metal ion and the amount of e repulsion between the ligand and the d e depends on the number of d e and hence the oxidation state of the metal.
Periodicity
the chemical and physical properties of elements repeat periodically
Non metal oxides responsible for acid rain
sulfur trioxide
SO3 + H2O -> H2SO4
sulfur dioxide
SO2 + H2O -> H2SO3
dichlorine heptoxide
Cl2O7 + H2O -> 2HClO4
dichlorine monoxide
Cl2O + H2O -> 2HClO
Ligand
species with lone pair of electrons which forms coordinate bond to metal ion (in complex)
