Topic 3: Periodicity Flashcards
atomic radius
the distance from the centre of the nucleus to the outermost shell of electrons
bonding atomic radius
- aka covalent radius
- half of the distance between the centre of the nuclei of 2 covalently-bonded atoms
why is atomic radius (RnB) > covalent radius (Rb)?
- the covalent bond is formed by the overlapping of atomic orbitals
- the overlapping region becomes common ground
periodicity
the tendency of similar properties of the element to recur at intervals, if they are arranged in order of increasing atomic number
periodic trend of atomic radius
- across a period, atomic radius decreases due to increased nuclear charge, which pulls electrons closer to the nucleus
- down a group, atomic radius increases due to increased no. of shells and increased shielding effect
ionic radius
the effective distance from the nucleus of the ion up to which it has an influence in the ionic bond
characteristics of ionic radius
- radius of cation radius of parent atom
electron affinity
- the amount of energy released when a mole of electrons is added to a mole of gaseous atoms
- represents the ability of an atom to hold an additional electron
- the higher the tendency, the higher the e.a.
factors affecting electron affinity
- nuclear charge: higher nuclear charge = higher e.a.
- size of atom: the larger the atom, the lower the e.a.
- e. config.: stabler configs = lower e.a.
periodic trends of electron affinity
- along a period, e.a. increases, because nuclear charge increases and atomic size decreases
- down a group, e.a. decreases, because although size and nuclear charge both increase, effect of increased size is much more pronounced
electronegativity
a measure of the ability of atoms to attract a shared pair of electrons toward itself in a covalent bond
polar covalent bond
bond in which electrons are shared unequally, resting in a partial polar charge
periodic trends of electronegativity
- along a period, electronegativity increases due to increased nuclear charge, resulting in increased attraction between nucleus and electrons
- down a group, electronegativity decreases due to increase in atomic size, resulting in reduced attraction
types of oxides
- acidic oxide
- basic oxide
- neutral oxide
- amphoteric oxide
Group I trends
Down the group:
- reactivity increases
- m.pt/b.pt increases
- density decreases
Group VII trends
down the group:
- m.pt/b.pt increases
- density increases
- reactivity decreases
transition element
- element whose atom has an incomplete d sub-shell
- or can give rise to a cation with an incomplete d sub-shell
properties of transition metals
PHYSICAL PROPERTIES
- high m.pt/b.pt
- high tensile strength
- high density
- malleable and ductile
- high thermal & electrical conductivity
CHEMICAL PROPERTIES
- form compounds with more than one oxidation no
- incomplete d-subshell
- form complex ions
- form colored compounds
- act as catalysts when either elements or compounds
why does chromium show a large jump in IE for the 7th e-?
Chromium’s config: [Ar] 3d5
- because there’s not much diff between 3d and 4s orbitals
- but taking from 3p6 is harder (also because the orbital is full)
first ionisation energy
- the amount of energy required
- to remove 1 mol of e-s from 1 mol of atoms in gaseous state
what does group number tell you about electron placement?
number of outer shell/valence electrons
what does period number tell you about electron placement?
number of occupied electron shells
Why do halogen m.pts/b.pts increase down the group?
- coz their Mr also increases down the group
- thus van der waals/intermolecular forces also increase
why do alkali metal m.pts/b.pts decrease down the group?
- metallic bonding weakens
- as radius increases, delocalized electrons are shielded more from the effective nuclear charge through electron shielding effect
why do halogen b.pts increase as Mr increases?
- coz their intermolecular attraction increases
- due to temporarily induced dipoles
explain why there’s an increase in m.pt from one side of a period to the other (cations only)
- cations become more positive
- size and radius decreases, so charge density increases
explain why there’s an increase in m.pt from one side of a period to the other (anions only)
- Mr increases
- van der waals/london forces also increase
why does Si have a greater m.pt than Ar?
- Si has a giant 3-D macromolecular covalent bonding structure
- while Ar is bonded by atomic van der waals/london/dispersion forces
is Zn a transition metal?
- it’s a d-block element but not a transition metal
- zinc has a complete d subshell in both neutral atom and ion form
- it also only has 1 oxidation state
is Sc a transition metal?
- it has no d electrons
- so it’s colorless in aqueous solutions
- but its atom has an incomplete d subshell
- and it has 2 oxidation states
- and the Sc2+ ion has a d e-
- Sc is a transition metal
why do transition metals have variable oxidation numbers?
- due to patterns in successive ionization energies
- outer d and s subshells in transition metals are closer in energy compared to other elements
- so they don’t experience the characteristic jump in IE from an s e- to a d e-
characteristic oxidation states shared by transition metals
- oxidation states correspond to the use of 3d and 4s e-s in bonding
- all transition metals can form ions of +2 and +3 oxidation states
- M3+ is the most stable oxidation state for Sc to Cr, but all other transition metals prefer M2+ due to increased nuclear charge making IE3 comparatively higher
highest possible oxidation state of a transition metal
+7 at Mn
can transition metals show covalent characteristics?
- yes
- at oxidation states > +3
- this is due to ions with higher charge having such a large charge density that they polarize cations and increase the covalent nature of the compounds
uses of compounds with high oxidation states
oxidizing agent
e.g. K2Cr2O7
complex ion
- transition metal ions in aqueous solutions that form coordinate bonds with water molecules
- due to the metal ions’ high charge density
ligands
- the species surrounding a central ion in a complex ion
- all ligands have at least 1 atom with a lone pair of e-s
how is a complex ion formed?
- forms when a central ion is surrounded by molecules/ions (ligands) with a lone pair of e-s
- the ligands attach via coordinate bonding
coordination number
the number of coordinate bonds formed by the central ion
polydentate ligand
- ligands that form more than 1 coordinate bond
e. g. as a hexadentate ligand, EDTA4- can form 6 coordinate bonds, and thus is equal to 6 monodentate ligands
chelate
compound containing a ligand that has occupied more than one bonding site (polydentate ligand)
importance of chelates
- can remove transition metal ions from solutions
- thus inhibiting enzyme catalyzed oxidation reactions
heterogenous catalysts
the catalyst is in a different state from the reactants
why do transition metals make effective heterogenous catalysts?
- they can use their 3d and 4s e-s to form weak bonds to reactants
- thus enabling them to collide with the correct orientation
examples of transition metals as heterogeneous catalysts
- Fe in Haber’s Process
- Ni in alkenes –> alkanes
- Pd and Pt in catalytic converters
- MnO2 in the decomposition of hydrogen peroxide
- V2O5 in Contact Process
homogeneous catalysts
catalyst in same state as reactants
why do transition metals make effective homogeneous catalysts?
- transition metal ions have variable oxidation states
- thus they are particularly effective catalysts in redox reactions
examples of transition metals as homogeneous catalysts
- Fe2+ in heme (haemoglobin)
- Co3+ in Vit B: 5 of the 6 sites are occupied by N, and the 6th site is used for biological activity
diamagnetism
- common to all materials
- as the orbital motion of electrons is what produces the diamagnetic effect
- it produces weak opposition to applied magnetic field
paramagnetism
- common to all substances with unpaired e-s (includes all transition metals and their compounds)
- stronger than diamagnetism
- produces magnetization proportional to applied magnetic field and in same direction
ferromagnetism
- occurs in substances with long range ordering of unpaired e-s
- strongest form of magnetism
- produces magnetization greater than applied magnetic field
e.g. iron, nickel, cobalt
domains
- occurs in ferromagnetism
- domains are regions in which unpaired d e-s in a large no of atoms line up with parallel spins
- they are randomly oriented but can be more ordered when a magnetic field is applied
- the magnetism will remain even after the magnetic field is removed
why do transition metals appear colored?
- transition metal ions absorb specific colors from the color spectrum
- in the presence of a ligand’s lone pair of e-s, their d orbitals split into 2 sub-levels
- upon absorbing light, a 3d e- becomes excited and passes into the higher energy level
factors affecting color of complex
- nuclear charge and identity of central ion
- charge density of ligand
- geometry of complex ion
- no. of d e-s present (i.e. oxidation no of central ion)
factors affecting color of complex: nuclear charge and identity of central ion
- the strength of the coordinate bond depends on the electrostatic attraction between the nuclear charge and lone pair of e-s
- higher nuclear charge = higher electrostatic attraction = higher energy wavelengths (i.e. shorter wavelengths) absorbed
factors affecting color of complex: charge density of ligand
- ligands with higher charge densities cause greater splitting of d orbitals
- greater splitting = greater energy of light (i.e. shorter wavelengths) absorbed
factors affecting color of complex: geometry of complex ion
the splitting in energy of d orbitals depends on the relative orientations of the ligands and d orbitals
factors affecting color of complex: no of d orbitals and oxidation state of central metal ion
the no of d orbitals and the oxidation state affect:
- strength of interaction
- amount of e- repulsion
between the ligands and the central ion
