PERIODICITY

Atomic radius

ionization energy

electron affinity

electronegativity

the distance from the valence shell to the nucleus

increases across the period, decreases down the group

the minimum amount of energy required to remove one elctron from one mole of a gaseous atom

increases across the period, decreases down the group

the change in energy when one electron is added to one mole of a gaseous atom

atomic radius
ENC

as the radius decreases (across the period), the electron affinity increases because we’re adding an electron to a positively charged ion.

as the ENC increases (across the period), the electron affinity increases due to the increase in ENC across the period, this means that the attractive forces increase as well.

an atom’s ability to attract a bonding pair of electrons

non-polar covalent bond (equal distribution)

polar covalent bond (unequal distribution, partially negative and partially positive).

very reactive - LOW IE

weak metallic bonding

• low MP / BP

• can be easily cut with a knie

low density

form white-colored compounds unless they react with a transition element (transition elements produce colored compounds, except zinc)

metallic

lattice rows of cations surrounded by a sea of delocalized electrons

radius increases, therefore charge density decreases = weaker attraction between cations and delocalized electrons

because they have the same number of valence electrons

reactivity increases down the group.

IE decreases down the group due to increase in radius size (number of energy levls increases), so, weaker attraction between the electron (valence) and the nucleus; easier to lose.

metal oxides

basic, non-metal oxides are acidic (ex: CO2)

metal hydroxides and H2 gas

2Na(s) + 2H2O(l) → 2Na+OH-(aq) + H2(g)

They are found as diatomic molecules in the element state (X2).

They are coloured:

• F2 yellowish gas

• Cl2 greenish gas

• Br2 deep red/ orange brown vapour.

• I2 dark grey solid/ purple fumes.

They have low melting and boiling points, that increase down the group. 4)

They have (7) valence electrons.

theyre non-polar, and are bonded by LDF of attraction. LDF of attraction depends on molar mass, which increases down the group. hence, when molar mass increases strength of LDF of attraction increases as well.

up the group, electron affinity increases as an electron is being added to an energy level closer to the nucleus, so more energy is released.

their salt, metal halides

ex: NaCl

metal gets oxidized, halogen gets reduced

down the group (IE decreases)

up the group, electron affinity increases and an electron is being added to an energy level closer to the nucleus.

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metallic

Na2O

basic

metallic

MgO

basic

metallic

Al2O3

amphoteric (acts as an acid or base)

giant covalent

SiO2

acidic

simple covalent

P4O6 // P4O10

acidic

simple covalent

SO2 // SO3

acidic

simple covalent

Cl2O // Cl2O7

acidic

LDF of attraction

no oxides formed

N/A

Na2O(S) + H2O(l) → 2NaOH(aq)

MgO(S) + H2O(l) → Mg(OH)2(aq)

P4O6(s) + 6H2O(l) → 4H3PO3(aq)

P4O10(s) + 6H2O(l) → 4H3PO4(aq)

SO2(g) + H2O(l) → H2SO3(aq)

SO3(g) + H2O(l) → H2SO4(aq)

nitrogen will only react with oxygen gas under the following severe conditions:

high temperature

high pressure

N2(g) + O2(g) → 2NO(g)

The nitrogen (II) oxide (nitrogen monoxide) reacts with Oxygen gas in the

atmosphere to form nitrogen (IV) oxide (nitrogen dioxide):

2NO(g)+ O2(g) → 2NO2(g) acidic oxide.
The nitrogen (IV) oxide dissolves and reacts in water to form an acidic solution.

NO N2O

2NO2(g) + H2O(l) → HNO3(ag) + HNO2(aq)

NO // N2O are neutral oxides.

elements with partially fileld d-orbitals

zinc, it has a partially filled d-orbital

high MP / BP / density

more than one oxidation state

form complex ions

form colored compounds

can be catalysts

• ex: Fe in haber process

exhibit magnetic properties

• paramagnetic

• ferromagnetic

high electrical / thermal conductivity

malleable and ductile

a type of magnetic behaviour that is caused by the presence of unpaired electrons.

IT IS AFFECTED BY THE MAGNETIC FIELD

The greater the number of unpaired electrons the greater the paramagnetic behaviour it displays; and vice versa.

a type of magnetic behaviour that is caused by the presence of paired electrons.

IT IS NOT AFFECTED BY THE MAGNETIC FIELD

electron is removed from s-orbital before d-orbital

(+2), because the electrons are removed from (4s) first.

Because their successive ionization energies are very close to each other.

complex ions

either a negative ion or a neutral molecule with a lone pair(s) of electrons

monodentate and polydentate

can donate one pair of electrons to be shared with the metal cation

ex: I-, OH-, SCN-, H2O, NH3, etc

can donate two or more pairs of electrons to be shared with the metal cation

ex: ethanedioate

using lone pairs

a dative bond

it is a covalent bond formed when a particle provides a lone pair of electrons to a particle that is deficient in electrons to be shared with the two atoms.

electron pair acceptor

METAL CATION

electron pair donor

LIGAND

the number of coordinate covalent bonds (dative bonds) formed between the ligand and the metal cation

catalysts

because of their variable oxidation state

Nickel as a catalyst for the hydrogenation of an alkene (into an alkane).

the amount of the splitting of the d-orbitals and the difference between the d-orbitals in the (d) sub-energy level.

if the light absorbed has a high frequency and short wavelength, then the emitted colour will have the opposite properties.

identity of metal

oxidation state of metal ion

identity of the ligand

the geometry of the complex ion

a. some metal ions with different oxidation states have different colors due to different electron configuration

IF THEY ARE ISOELECTRONIC / NUCLEAR CHARGE

As the nuclear charge increases; the ligands are pulled more towards the metal cation, causing more splitting in the d-sub-energy levels; light with high frequency and short wavelength will be absorbed but the complementary color will have lower frequency and longer wavelength and vice versa.

Ions of the same metal with different oxidation numbers will have different colours.

ex:

[Fe (H2O)6]2+ Pale green

[Fe (H2O)6]3+ Orange

That can be explained for two reasons

• different electron configurations.

• ions with a higher oxidation state will have a higher ability to attract the ligands towards the nucleus = greater splitting of d-orbitals

in accordance to the spectrochemical series (found in data booklet)

I− < Br− < Cl− < F– < OH− < H2O < NH3 < CO ≈ CN−

the wavelength of visible light

d-d transition