PERIODICTY Flashcards
Are noble gases hard to remove electrons and why?
Yes, full electron configuration
explain the trend in reactivity of the halogens
reactivity decreases down the group, as atomic radius increase (more electron shells), attraction of nucleus on electron decreases (electron affinity decreases)
describe the bonding in metals and explain their malleability
electrostatic attraction between a lattice of positive ions/cations and delocalised/sea of electrons
3 characteristic properties of transition elements
variable oxidation numbers/valency
if ions they form complex
form coloured COMPOUNDS/IONS
catalytic behaviour
what is the easiest group to remove electrons from?
Group 1
what do oxidising agents do?
Gain electrons (reduced)
the oxide that metal creates, are they acidic or basic?
Basic
what must all ligands contain?
a non-bonding pair of electrons
what is “Electron affinity”
The energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.
ionisation energy (IE)
is the amount of energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.
draw out the period table relationships
this includes:
atomic radius
ionisation energy
electronegativity
electron affinity
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Properties of non-metals and metals
include:
density, ionisation energy
electronegativity values
cations or anions
reducing and oxidising
what types of bonds with non-metals
basic or acidic oxides
metals:
high density
low ionisation energies
low electronegativity values
lose electrons to form positive cations
behave as reducing agents (undergoes oxidation)
form ionic bonds with non-metals
form basic oxides
non-metals:
low density
high ionisation energies
high electronegativity values
gain electrons to form negative ions (anions)
behave as oxidising agents (undergo reduction)
form covalent bonds with other non-metals
form acidic oxides
properties of metalloids
moderate density, intermediate ionisation energies, intermediate electrical conductivity, amphoteric or weakly acidic oxides
what is effective nuclear charge
is the net positive charge experienced by valence electrons
Z(eff) = Z atomic number - S number of shielding electrons
does Z eff decrease or increase across a period (left to right)
increase
what does Z eff do when doing down a group
Stays the same
what is electron negativity
Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond.
size of cation, atom and anion compare
cation < atom < anion
outline 2 reasons why electronegativity increases across period 3 in the periodic table and 1 reason why noble gases are not assigned electronegativity numbers
- because atomic radius decreases
- nuclear charge increases (number of protons)
- Nobel gases have full valence electron shell
What causes transition metals to have colour?
electron transition between d-orbitials
the ___ the ligand is in the electrochemical series, the ____ the splitting of the d orbitals
the HIGHER the ligand is in the electrochemical series, the GREATER the splitting of the d orbitals
Left to right on the periodic table, does it go basic to acidic or acidic to basic?
basic to acidic (think of H2SO4 as it is acidic)
what is periodicity?
repeating patterns of physical and chemical properties
describe the structure of a complex ion
A complex ion consist of a central metal ion bonded to ligands by coordinate covalent bonds
define the term ligands
a ligand is a species with a lone pair of electrons that forms an coordinate bond to the central central metal ion. Ligands can be negative
ions (CN-1) or neutral molecules (NH3).
how does a coordinate covalent bond differ from a conventional covalent bond
In a coordinate covalent bond, both electrons in the bond comes
from one atom. In a conventional covalent bond, each atom
supplies one electron in the bond.
explain how ligands are also able to act as a lewis base
Ligands have lone pairs of electrons, therefore they can act as
Lewis bases. The central metal ion is classed as the Lewis acid.
what is diamagnetic
when you have no unpaired electrons in the d orbitals
what is paramagnetic
when you have unpaired electrons in the d orbitals
what is the charge on ligand:
H2O
NH3
CN
Cl
SO4
H2O = 0
NH3 = 0
CN = -1
Cl = -1
SO4 = -2
4 ligands is…
Square planar or tetrahedral
can cations be ligands
no
what is the product of Cl2 + H2O
HCl + HOCl
the halogen at the ___ of the group _____ the halide ____
the halogen at the top of the group displaces the halide below
a higher electronegativity correlates with higher… and what does it usually involve?
boiling and melting points, intermolecular forces
Suggest why the colour of [Fe(NH3)6]3+ is different from the colour of [Fe(H2O)6]3+
the only diff is the ligand. NH3+ has a higher charge density than H2O so as it approaches the ion it creates a larger split in the d-orbital electrons and a lower wavelength is absorbed
Suggest why the colour of [Fe(NH3)6]2+ is different from the colour of [Fe(H2O)6]3+
The oxidation number. a different number of d electrons. Fe 3+ has the higher oxidation number so shorter wavelength is absorbs and a larger energy difference
suggest why the colour of [Cr(H2O)6]3+ is different from the colour of [Fe(H2O)6]3+
the only diff is the central metal ion thus Fe has a higher nuclear charge than Cr/ Ligands interact more effectively with the d-orbitals of an ion with higher nuclear charge thus a shorter wavelength is absorbed and a larger energy diff between Fe and Cr.
Explain why Fe2+(aq) is coloured and can behave as a reducing agent, whereas Zn2+(aq) is not coloured and does not behave as a reducing agent.
Zn2+ ion has a full 3d orbital, is not a coloured transition metal. Fe2+ does not have full 3d orbitals.
what does charge density influence?
spin and creates a larger spilt. The bigger the charge density the bigger the split and spin.
