cosmic abundance of the elements - chemistry coursescourses.chem.psu.edu/chem112/fall/lecture...

1 LRSVDS Chem 112 Main Group Elements

Descriptive Chemistry

Cosmic abundance of the Elements

Element Atomic No. Rel. Abundance

H 1 3.5 x 108 He 2 3.5 x 107

C 6 80,000

N 7 160,000

O 8 210,000 F 9 90

Ne 10 10,000

Na 11 462

Mg 12 8,870

Al 13 882 Si 14 10,000

Cr 24 95 Mn 25 77

Fe 26 18,300


Occurrence of the Elements on Earth


Abundance of elements

Terrestrial abundance; On Earth!

Element Atomic No. PPM Rank

H 1 8,700 9 He 2 0.003

C 6 800 N 7 300

O 8 495,000 1

F 9 270

Na 11 26,000 6 Mg 12 19,000 8

Al 13 75,000 3

Si 14 257,000 2

K 19 24,000 7 Ca 20 34,000 5

Ti 22 5,800 10

Fe 26 47,000 4


Abundance of elements Terrestrial abundance

0

1

2

3

4

5

6

7

8

9

10

H B C N F Na Mg Al P O S Cl K Ca Ti V Cr Mn Fe Co Cu Zn Si Se Rb Mo Sn I

% b

y w

eigh

t

Si 27.72% O 46.6%


Isolation of the Elements From their Natural state

Sources of Gas-Phase Nonmetals:

1) HEAVY Noble Gases, O2, N2

2) Helium

3) Halogens

a)

b)

4) Hydrogen

a) Steam Reforming

CH4(g) + H2O(g) ! CO(g) + 3 H2(g)

CO(g) + H2O(g) ! CO2(g) + H2(g)

b)

c)


Isolation of the Elements From their Natural state

Sources of Other Nonmetals: 1) Carbon

2) Sulfur

3) P

4) Se, Te

5) ! Boron

6) ! Metalloids


Periodic Trends Summary

Zeff

n

1) Atomic Size 2) ionization energy

3) electron affinity

both + and - ; halogens most negative

4) metallic character

- metals lose electrons

- nonmetals gain electrons

increasing metallic character


N Max. of 4 bonds, 1 lone pair; usually has a valence of 3

HNO3 (4 bonds to N)

NH3, NCl3, CH3NH2, HNO2

O Max. of 4 bonds, 2 lone pairs; usually has a valence of 2

H2O, OF2, H2C=O

C Max. of 4 bonds, 0 lone pairs; always has a valence of 4

can gain 4 or lose 4 electrons to make an octet

So carbon always makes 4 bonds

CH4 (4 single bonds)

O=C=O (2 double bonds)

H-C!C-H (1 single + 1 triple bond)

H2N

C=O (2 single + 1 double bond)

H2N (urea)

Valence in Period II


Bonding in Period II vs. III Valence: Heavy Group V elements can gain 3 electrons or lose 5 to make an

octet; " expect either 5 covalent bonds (PF5) or 3 covalent bonds (PF3)

Multiple Bonds: elements past the second row are too big to allow good sidewise overlap of p-orbitals


Difference in Bonding: Period 2 vs. 3

O2 is molecular (O=O, has a double bond)

But S forms rings with single bonds (e.g., S8)

N2 has a triple bond (:N!N:, very stable)

But phosphorus is found in several forms (white, red, black),

all of which have only single bonds.


Examples: HCl, HBr, CH4, H2Se, SiH4, …

acidic or show no acid-base properties

acid strength increases from left to right PH3 < H2S < HCl

acid strength increases going down family H2O < H2S < H2Se < H2Te

NH3 and other amines are basic (lone pair)

Chemistry of Hydrogen Forms binary hydrides (hydrogen and one other element)

ionic, metallic or molecular hydrides

Molecular hydrides: H bonded covalently to a non-metal


Examples: LiH, LiAlH4, CaH2

these metals are much less electronegative than hydrogen;

H gains electrons, produces a

hydride anion H -

**high melting ionic solids, very basic, strong reducing agents

CaH2(s) + H2O (l) # H2(g) + Ca(OH)2 (aq)

Metallic hydrides: H plus a transition metal many retain their metallic characteristics

Ionic Hydrides: Hydrogen plus a group 1 or 2 metal

If ratio of M:H is not fixed;

hydrogen atoms are absorbed into the interstices of the metal

lattice

If M-H ratio is fixed;

organometallic metal hydrides or metal-

dihydrogen complexes Fe(H2)(H)2 (PR3) 3


"!Only group 3 element that is a nonmetal

"!Network covalent solid (mp=2300ºC)

Sources of Boron: Rare, found in one mineral Borax: Na2B4O7

.(OH2)10

Chemistry of Boron:

"!Oxide B2O3 is used to make Pyrex glass (anhydride of boric acid H3BO3)

"!Boranes (B plus H) are

electron deficient (6 v.e.)

"!Borohydrides (borane anions); source of H! ions

"!Boron Nitride, BN

Boron is one element to the

left of carbon. Nitrogen is one element to the

right of carbon.

Properties of Boron Group III; B, Al, Ga, In, Tl

BN hardness = 9.8


Trends in Group III; B, Ga, In, Tl

•! Inert Pair Effect

•! On descending the group +1 oxidation state becomes more stable than +3

•! Oxides and hydroxides become more basic

Going down the group metallic character increases:

B(OH)3 is a weak acid

Al(OH)3 and Ga(OH)3 are amphoteric

In(OH)3 is basic

•! Bonding: covalent # ionic going down the group

B is most electronegative (ONLY covalent bonds)

Al, Ga, and In form both covalent and ionic bonds.

The Inert Pair Effect: As you get closer to the bottom of the group, there is an increasing

tendency for the s2 electrons not to be used in bonding.


Group IV: C, Si, Ge, Sn, Pb Chemistry of carbon

#!Organic Chemistry (C bonded to H)

#!Inorganic Chemistry (C not bonded to H)

$! Oxides CO, CO2

$! Carbides

1. ionic (e.g., CaC, C with active metals )

contain C4- or C22 - (-C!C-)

C4- : Be2C, Al4C3

reacts with water to form CH4

C22 - : CaC2

reacts with water to give HC!CH

2. covalent (e.g., SiC, covalent bonds with metalloids and nonmetals)

SiC: does not react with water; very hard

3. interstitial (e.g., steel, C incorporated into interstitial spaces of transition metals in non-stoichiometric proportions)

•!steel is an interstitial carbide: harder than pure Fe


TRENDS IN GROUP IV Going down the periodic table in Group IV:

1)! The +2 oxidation state becomes more stable than +4 due to the “inert pair”

effect. +2 is rare for C, Si, Ge.

+2, +4 is common for Sn.

+4 is unstable for Pb " strong oxidizing agent

" prefers to be +2.

2)! Basicity of oxides and hydroxides increases

CO2, SiO2, GeO2 are weakly acidic. SnO, SnO2, PbO are amphoteric.

3)! Hydrides become less stable

THOUSANDS of stable hydrocarbons (compounds of C and H)

SiH4 is stable but is spontaneously flammable. Ge, Sn, Pb hydrides are very unstable.


Group 5-6: Hydrolysis of nitrogen oxides Hydrolysis = reaction with water

N is a non metal: oxides are acidic.

N-oxide + H2O = oxyacid

N2O5 + H2O # 2HNO3 (nitric acid)

3NO2 + H2O # 2HNO3 + NO

N2O3 + H2O # 2HNO2 (nitrous acid)

HNO3 is colorless, corrosive; turns yellow

in sunlight (photochemical decomposition)


Trends for GROUP V ELEMENTS;

N, P, As, Sb, Bi Trends Going down the periodic table:

1) ! Electronegativity decreases.

2) ! Size increases

3) Switch from non-metallic to metallic

4) Hydroxides and oxides start acidic, become more basic.

5) ! Hydrides become less stable.

NH3 is stable.

PH3 is stable but burns in air. AsH3 decomposes easily.

SbH3, BiH3 are very unstable.

7) ! “Inert pair effect” becomes more pronounced: +3 becomes more

stable than +5.

P: +5 dominates (lower EN than N)

As3+, As5+ are equally common. Sb: +3 dominates.

Bi: +3 dominates.


General Properties

Most widely used oxidizing agent O-O bond in O2 is strong (495 kJ/mol), as are many X-O bonds

ns2np4 configuration;

can gain 2 e- or share 2 e- to fill the shell

" Oxidation states

-2 (O in all compounds except with F; O2F2, OF2 ) -1 (only in peroxide ion, O2

2-)

+2, +4

+6 (NOT O, only larger members of group)

Group VI Elements: O, S, Se, Te (Po)

Allotropes of oxygen: Light or

electrical discharge

3O2 2 O3

decomposition

Strong oxidizing agent

O3 # O2 + O

O2: Most abundant element on earth (21% of air)

50% of earth’s crust is oxygen (but not O2)


Anhydrides and Dehydration Anhydride: compound formed by loss of ___________________

2 NaOH ! H2O + Na2O base anhydride

H2SO4 ! H2O + SO3 acid anhydride

H3PO4 ! H2O + H4P2O7 acid anhydride

Anhydrides are good

dehydrating agents;

They react with water (hydrolysis rxn)

Sucrose + H2SO4 product:

(C12H22O11) (removes H2O) carbon


Halogens (Group 7) •! Have valence of 1 (same as H)

Can be used to replace H in many compounds thus influencing the properties.

•! High reactivity of halogens WHY??

•! Can also have positive oxidation states when combined with more electronegative elements

(e.g. oxygen to form oxyhalides: ClO4–, IO3

–, etc.)

Selected uses

Fluorine – prepare polymers, fluorocarbons (Teflon,

CFC’s), glass etching

Chlorine – chlorinated organics/plastics (PVC), bleaching

agent (NaClO), water treatment (destroy bacteria)

Bromine – photographic film (AgBr)

Iodine – iodized salt


INTERHALOGEN COMPOUNDS: Combine halogens; central atom is least electronegative

(it has to share electrons with surrounding atoms)

1:1 ratio BrF, ClF, ICl, BrCl

1:3 ratio ClF3, BrF3, IF3 , ICl3 (the only non-F)

1:5 ratio BrF5, IF5 (ClF5 formed with difficulty)

Only Br & I large enough to be central atoms with 5 surrounding atoms

1:7 ratio IF7

Only I is large enough to be a central atom with 7 surrounding atoms

Central atom is always bigger in size and less electronegative

Extremely Reactive: superb oxidizing agents

Oxidation state of central atom decreases to 0 or -1


Chemistry of Noble Gases; Group 8 React only under rigorous conditions

Heavier elements have lower ionization energy, can more easily share e-

Xe Ionization E =1176 kJ

O2 Ionization E = 1171kJ

"!Xe forms compounds with F and O!

Xe + n F2 # XeF2

OR XeF4

OR XeF6

Oxidation states of +2, +4, +6

XeF6(s) + H2O(l) # XeOF4(l) + 2HF(g)

XeF6(s) + 3H2O(l) # XeO3(l) + 6HF(aq)

Kr forms only one compound; KrF2 Only one known “compound” made with He; Endohedral complex of He

inside a C60 cage

cosmic abundance of the elements - chemistry coursescourses.chem.psu.edu/chem112/fall/lecture...

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