PERIODICITY

1-56 SL MATERIAL 57 - x HL material

Physical properties

Atomic radius

ionization energy

electron affinity

electronegativity

atomic radius, definition

the distance from the valence shell to the nucleus

atomic radius, trend

increases across the period, decreases down the group

ionization energy, definition

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

ionization energy, trends

increases across the period, decreases down the group

electron affinity, definition

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

factors affecting electron affinity

atomic radius
ENC

factors affecting electron affinity: radius

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

factors affecting electron affinity, ENC

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.

electronegativity

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

results of adding a bonding pair of electrons

non-polar covalent bond (equal distribution)

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

properties of alkali metals; elements of group 1

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)

properties of alkali metals; elements of group 1

type of bonding

metallic

what is metallic bonding

lattice rows of cations surrounded by a sea of delocalized electrons

melting point of alkali metals decreases down the group

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

alkali metals have similar chemical reactivity

because they have the same number of valence electrons

reactivity trend

reactivity increases down the group.

why does reactivity increase 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.

alkali metals and oxygen form…?

metal oxides

pH of metal oxides (acidic or basic)

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

alkali metals and water form…?

metal hydroxides and H2 gas

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

Properties of the halogens

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.

BP of halogens increases down the group

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.

reactivity of halogens increases up the group (opposite to alkali metals)

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.

halogen X alkali metal, what do they form?

their salt, metal halides

ex: NaCl

which gets oxidized and which gets reduced?

metal gets oxidized, halogen gets reduced

trend in increasing reactivity: metals

down the group (IE decreases)

trend in increasing reactivity: halogens

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

the oxides of period 3 (Na, Mg, Al, Si, P, S, Cl, Ar)

next slide

Na: Bonding

metallic

Na: oxide formes

Na2O

Na: Nature of oxide

basic

Mg: Bonding

metallic

Mg: oxide formed

MgO

Mg: nature of oxide

basic

Al: Bonding

metallic

Al: oxide formed

Al2O3

Al: nature of oxide

amphoteric (acts as an acid or base)

Si: Bonding

giant covalent

Si: oxide formed

SiO2

Si: nature of oxide

acidic

P: Bonding

simple covalent

P: oxide formed

P4O6 // P4O10

P: nature of oxide

acidic

S: Bonding

simple covalent

S: oxide formed

SO2 // SO3

S: nature of oxide

acidic

Cl: Bonding

simple covalent

Cl: oxide formed

Cl2O // Cl2O7

Cl: nature of oxide

acidic

Ar: Bonding

LDF of attraction

Ar: oxide formed

no oxides formed

Ar: nature of oxide

N/A

Equations that shows the basic behaviour of the oxides of the elements of period 3.

Na2O(S) + H2O(l) → 2NaOH(aq)

MgO(S) + H2O(l) → Mg(OH)2(aq)

Equations that shows the acidic behaviour of the oxides of the elements of period 3.

P4O6(s) + 6H2O(l) → 4H3PO3(aq)

P4O10(s) + 6H2O(l) → 4H3PO4(aq)

SO2(g) + H2O(l) → H2SO3(aq)

SO3(g) + H2O(l) → H2SO4(aq)

nirtogen oxides, conditions of formation

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.

what are transition elements?

elements with partially fileld d-orbitals

exception to transition elements

zinc, it has a partially filled d-orbital

Properties of transition elements

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

define paramagnetic

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.

define diamagnetic

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

IT IS NOT AFFECTED BY THE MAGNETIC FIELD

ionization of transition elements

electron is removed from s-orbital before d-orbital

what is the most common oxidation state for the transition elements and why?

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

why do transition elements have variable oxidation states

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

transition elements form…

complex ions

define ligand

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

what are the two types of ligands

monodentate and polydentate

monodentate

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

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

polydentate

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

ex: ethanedioate

how does a ligan bond to a metal cation so it can form a complex ion?

using lone pairs

what kind of bond is formed between a ligand and the metal cation?

a dative bond

define 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.

Important definition: Lewis acid

electron pair acceptor

METAL CATION

Important definition: Lewis base

electron pair donor

LIGAND

define coordination number

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

what can transition elements be used as?

catalysts

why can transition elements be used as catalysts?

because of their variable oxidation state

example of a transition element as a catalyst

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

Factors that affects the color of transition metal complexes

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.

factors affecting the ammount of splitting of a d-orbital

identity of metal

oxidation state of metal ion

identity of the ligand

the geometry of the complex ion

identity of the metal

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.

oxidation number

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

nature of the ligand

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

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

what does the energy gap between the two sets of d-orbitals represent?

the wavelength of visible light

the transition of electrons from the low-energy d-orbital to the high energy d-orbital is called…?

d-d transition