General trends in main group elements - … trends in main group elements: Physical properties: ... (Ra). - Alkaline earth metals have only two electrons in their outermost electron

General trends in main group elements: Physical properties:

- They complete their Electronic Configuration using S and P electron subshells.

- Group number indicates the number of electrons in the outermost shell (valence electrons) of an atom.

- Far left alkali and alkaline earth metals with the expected metallic characters like luster, high ability to conduct heat and electricity (used to distinguish between metals and nonmetals) and that is because they contain loosely bound valence electrons that are free to move and therefore conduct the current.

- Far right are nonmetals with localized or covalently bonded electron pairs and do not conduct the current.

- Elements diagonally from Boron (B) to polonium (Po) are intermediate having both metallic and nonmetallic characters (semi-metals).

Introduction

Electronegativity

- Is the ability of the atom to attract electrons towards itself.

- Increases cross the period and decreases down the group following the size of the atom.

- Characteristic of nonmetals so fluorine with the highest electronegativity.

- Noble gases matches or exceed halogens because they are smaller in size.

Ionization Energy:

- Is the energy needed to completely remove the outer most electron/s.

- Follow the same trend as electronegativity.

- Decrease down the group and increase across the period with some exceptions for example :

I.E for B is less than Be and O is less than N even though the size is smaller.

This is because Be 2S2 full subshell and N 2P3 half full (very stable) While:

B 2P1 have outer electron in a higher energy level and can be easily lost.

Also O 2P2 the second electron must pair with the other e- which resulted in electron-electron repulsion that makes the lost of an electron very easy.

main group elementsof the Electronegativity

for the main group elementsIonization energy

H1Hydrogen:

- Electronic configuration: 1S1

- Position in the periodic table; can be placed:

1/ with gp 1A (Alkali metal) because:

a/ it is electro positive,

b/ with one electron in its outermost shell,

c/ valency = 1.

2/ With gp 7A (Halogens) because:

a/ it is a gas,

b/need one electron to fill outer shell,

c/ form -1 ion (H-),

d/ forms diatomic molecule H2.

3/ With gp 4A above carbon because:

a/ it is half filled orbital,

b/ same electroneg.,

c/ forms covalent bond.

BUT

Hydrogen does not belong to any of these gps and it has a unique characters not found in any other element owing to its small size since it is the smallest atom in the p. table and deserve special consideration.

Occurrence: - One of the most abundant element, the third most abundant in earth.

- It occurs mainly in the form of water or combined with carbon in organic molecules.

Preparation:

Laboratory:

1/ Reaction of metals with water at room temperature:

2Na(s) + 2H2O → 2NaOH + H2(g)

Ca(s) + 2H2O → Ca(OH)2 + H2(g)

2/ Less reactive metals (Mg) displace hydrogen from water steam to produce H2:

Mg(s) + H2O(g) → MgO + H2(g)

3/ Hydrogen is prepared in the laboratory by the action of acids on metals:

Zn + 2H+ → H2 + Zn2+

Commercially :

1/ Water-gas shift

Is the reaction between coke and steam at 1000 °C to produce a gaseous mixture known as water gas used as a fuel:

C(s) + H2O(g) → CO(g) + H2(g)

To separate H2 from CO the water gas passed with steam over a catalyst at 500 °C to convert CO into CO2:

CO(g) + H2O(g) → CO2(g) + H2(g)

CO2 removed by scrubbing:

Ca(OH)2 + CO2 → CaCO3 + H2O

2/ Steam Reforming

Is the reaction of the natural gas with steam in presence of a Ni catalyst:

CH4 + H2O → CO + 3H2

3/ Cracking:

Breaking large petroleum hydrocarbons into smaller molecules using solid catalyst to produce H2 and alkene:

C2H6 → C2H4 + H2

Alkane Alkene

Hydrogen Isotopes: Hydrogen occurs as three isotopes:

1/ Protium : ordinary hydrogen 1H or P, containing one proton and one electron.

2/ Deuterium: 2H or D, containing one proton, one electron and one neutron.

3/ Tritium: 3H or T, containing one proton, one electron and two neutrons.

Radioactive: 1H (P) (99.98% ) the rest is 2H (D), where as 3H (T) is radioactive traces only on earth.

For every H compound prepared there is an equivalent D compound for example, HCl and DCl, NH3 and DH3 , H2O and D2O.

3H (T) is produced by bombardment of the Li nuclei with neutron to produce He and 3H (T):

𝑳𝒊 + 𝒏𝟎𝟏 → 𝑯𝒆𝟐

𝟒 + 𝑯𝟏𝟑

𝟑𝟔

Because of the increasing atomic weight , the melting points and boiling points increase in the order:

P < D < T

Chemical properties

The chemistry of hydrogen depends mainly on three electronic processes:

1/ Gain an electron:

To reach the stable noble gas configuration forming the hydride ion H-.

This is very reactive with metals (alkaline metals and the alkaline Earth metals) to form simple (salt-like) hydride:

2Na(s) + H2(g) → 2NaH(s)

These hydrides are ionic and react with water to form hydrogen.

NaH + H2O → H2(g) + NaOH

2/ Loose an electron:

To form the positive ion H+ (proton) with very small size so it is found associated with solvent molecules. In water is found in the form H3O+

(hydronium ion).

3/ Share electron:

Hydrogen also forms covalent hydrides with carbon and other nonmetals by sharing electrons.

Also with fluorine and chlorine at room temp. to form colorless, covalent gases:

H2(g) + X2(g) → 2HX(g) X= F or Cl

When X is Br or I (less reactive) the reaction require higher temp.

- Hydrogen react with N at high temp. and pressure to produce ammonia (Haber Process):

3H2(g) + N2(g) → 2NH3(g)

The bonding in all these compounds is polar covalent .

The Hydrogen bond: Is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule and has one or more lone pairs of electrons.

a/ In these compounds the electronegative element strongly attract the bonding electrons and the H then has the partial positive charge (Hᵟ+).

b/ Hydrogen bonding increases the boiling point.

For example, NH3, H2O and HF (the first members of gp 5, 6 ,and 7A) have higher b.pt compared to the other members of the series.

In each of these three compounds , H-bonding makes it difficult to separate the molecules from the liquid state.

- H-bonding occurs not only between similar molecules but also occurs between different molecules.

1/ The molecule supplying the proton for H-bonding must be strongly polar (high electronegative) so that the hydrogen becomes highly positive. So the H-bond increases in the order:

N-H….N < O-H…O < F-H….F

Because electronegativity Of the atom bonded to hydrogen increases in the order:

N < O < F 2/ The proton acceptor ( that supplies the two electron for the H-bond) must be small for the bond to be strong one. For example Cl has the same electronegativity as N but it forms a weak H-bonding compared to N , this is because Cl atom is larger than N atom.

Two requirements for strong H-bonding:

Grp 1 (IA) ( Alkali Metals)

- The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr). They is found in the first column of the periodic table.

- Atoms of the alkali metals have a single electron in their outermost level, in other words, 1 valence electron.

- They are the most reactive metals.

- They react violently with water.

- Alkali metals are never found as free elements in nature. They are always bonded with another element.

- Two of the alkali metals, sodium and potassium, are essential for human life.

- The alkali metals are highly reactive solids having low melting points. They are ordinarily stored under nonreactive oil to prevent air oxidation, soft enough to be easily cut with a knife.

- Their melting points decrease with increasing atomic number(down the group) because metallic bonding between the atoms becomes weaker with increasing atomic size.

Extraction: - Potassium and sodium were first isolated as products of the electrolysis of molten KOH and NaOH.

- The following element isolated is lithium, by electrolysis of molten Li2O.

- Cesium and rubidium discovered after, and they were named after the colors of their emission lines (Latin, caesius, sky blue, rubidus, deep red).

- Francium a short-lived radioactive isotope.

Physical properties of the alkali metals are summarized in the following table:

Chemical properties: Grp IA

Oxidation states: Very similar in their chemical properties:

- Easily lose one electron (lowest ionization energies ) to achieve a noble

gas configuration forming +1 ion.

- No other oxidation state is known.

- No second ionization energy.

- Chemical bond is mainly ionic but there are some degree of covalency in some cases (greater with Li and least with Cs because of the increasing size).

- All are highly reactive metals.

- Excellent reducing agents because they can easily looses electrons.

- The reactivity of alkali metals increases on going from Li to Cs.

- Reactivity towards water: the metals react vigorously with water to form hydrogen; for example:

2 M + 2 H2O → 2 MOH + H2

(M = an alkali metal)

The reactivity towards water (and other chemical reagents) increases with

increasing electropositive nature (down the group from Li to Cs).

- Reactivity towards oxygen: alkali metals react with oxygen to form oxides, peroxides, and superoxides, depending on the metal.

- Reactivity towards di-hydrogen: all alkali metals combine with hydrogen to form hydrides of the formula MH.

2 M(s) + H2(g) → 2MH(s)

Reactivity of alkali metals

- Reactivity towards air: combustion in air gives the products (see table below).

- All alkali metal hydrides are ionic solids with high melting points. - The ionic character of the hydrides increases from Li to Cs.

- Reactivity towards halogens: alkali metals combine with halogens directly to form metal halides.

2 M + X2 → 2 MX

- The alkali metals dissolve in liquid ammonia. The solutions are coloured and metastable. - Solutions of alkali metals in liquid ammonia believed to contain solvated electrons: Because of these solvated electrons they: a/ conduct electricity b/ are excellent reducing agents.

- Lithium halides are some what covalent. It is because of the high polarization capability of lithium ion. The Li+ ion is very small in size and has high tendency to distort electron cloud around the negative halide ion. Since anion with large size can be easily distorted, among halides, lithium iodide is the most covalent in nature.

Grp 2 (IIA): Alkaline earth metals

- The alkaline earth metals are all of the elements in the second column (column 2A) of the periodic table. This group includes beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra).

- Alkaline earth metals have only two electrons in their outermost electron layer.

- Compared to grp 1 elements, these are harder, have higher melting points and boiling points, and are less reactive.

Occurrence and Extraction: Because of their reactivity, the alkaline earth metals

are not found free in nature.

- All can be obtained by electrolysis of their fused chlorides.

- Magnesium and calcium are among the most abundant elements in the Earth's crust (dolomite, calcite… and many other minerals).

- Strontium and barium are less abundant; they occur as sulfates and carbonates.

Physical properties of the alkaline earths are given in the following table:

- Beryllium is fifth in abundance of the alkaline earths and is obtained primarily from the mineral beryl (Be3Al2(SiO3)6). - All isotopes of radium are radioactive; it was first isolated by Pierre and Marie Curie from the uranium ore .

:Oxidation state - With the exception of beryllium, they have very similar chemical properties.

- Loose two electrons to achieve a noble gas electron configuration and form M+2.

- Good reducing agents.

- Atoms of the Grp 2 (IIA) elements are smaller than the neighboring Grp 1 (IA) elements because of the greater nuclear charge of Grp 2. So they are: - more dense; - have higher ionization energies than the Group 1; - have higher melting and boiling points; - widely used in alloys with copper, nickel, and other metals. - Radium used in the treatment of cancer.

Chemical properties: Grp 2

- Reactivity towards water and air: metal reactivity with water and air increases with atomic weight.

- Beryllium and magnesium are kinetically inert to oxygen and water. - Calcium , strontium and barium readily react with water and air.

- Reactivity towards halogens: the alkaline earth metals readily react with halogens at elevated

temperatures to form the halides of the type , MX2: M + X2 MX2 (X = F, Cl, Br, I)

- Reactivity towards hydrogen: all the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2:

M + H2 MH2

- Reactivity towards acids: the alkaline earth metals readily react with acids liberating di-hydrogen:

M + 2H+ M2+ + H2

- Beryllium is : the smallest different from the other alkaline earths in its chemical properties. forms covalent rather than ionic bonding.

Reactivity of alkaline earth metals

Beryllium and its compounds are extremely toxic.

Beryllium halides BeX2 may be monomeric and linear in the gas phase at high temperature,

In the crystal (solid) the molecules polymerize to form halogen-bridged chains, with tetrahedral coordination around beryllium,(see below).

Beryllium hydride, BeH2, is also polymeric in the solid, with bridging hydrogen's.

One of the most chemically useful magnesium compounds are the:

Grignard Reagents:

Formula RMgX (R = alkyl or aryl).

Found consisting of many species in solution linked by equilibria as those shown in the figure below:

positions of these equilibria, and the concentrations of the species depends on:

a/ The nature of the R group.

b/ The halogen.

c/ The solvent, and

d/The temperature.

Grignard reagents can be used to synthesize a range of organic compounds, including alcohols, aldehydes, ketones, carboxylic acids, esters, and amines.