Lecture 03 Chem 2

Transcript Lecture 03 Chem 2

Chemical Properties of Minerals II Basic Coordination Chemistry

Quantum theory gives us insight into the electronic structure of atoms and allows us to rationalize the biological behavior of minerals

Why we need to know the principles of chemistry in a minerals course: Minerals are chemicals that function in a biological setting.

Minerals perform functions that are attuned to their chemical properties Minerals are ions whose charge is determined by chemical principles Recognizing that most minerals exist as complexes with proteins and other organic molecules, their chemistry gives us insight into how these interactions take place In constructing life nature drew from a large pool of chemicals in an attempt to find the ones that best fit the tasks that had to be performed. Chemistry tells us how these decisions were made.

Electronic Structure is Behind Each of the Following Questions

1. Why do Ca 2+ , Mg 2+ , and Zn 2+ exist only as +2 ions? Li + , Na + and K + as +1 ions? 2. Why are bio-complexes of iron red and potassium and sodium colorless?

3. Why are zinc complexes with proteins stable while sodium complexes fall apart?

4. Why was calcium chosen to become the crystalline component of bone?

5. Why is zinc able to block the absorption of copper in the intestine?

6. Why is arterial blood cherry red while venous blood is a darker red?

7. What makes carbon monoxide gas so deadly?

8. Why are plants green?

9. Why is a dangerous oxygen radical formed when iron reacts with hydrogen peroxide?

10. Will the same happen with zinc and hydrogen peroxide?

The Basics

Insights into the Electronic Structure of Atoms

Energy is being emitted discontinuously

White lt Ca Ba

20 56 Fe Li 26 3 Emission Spectra of Elements

Pauli: electron exists in two different states Intrinsic electron spin

Conclusions:

Electrons are arranged in a very specified manner around the nucleus of atoms Quantum theory: “the electronic energy of atoms are quantizied….meaning they can only take on certain discrete energy values”.

A direct indication of the arrangement of electrons around a nucleus is ionization energies…the energy required to remove an electron from a gaseous neutral atom.

Some electrons are very labile Some electronic states are stable

1926 Erwin Schr ödinger likened the motion of electrons around a central nucleus of an atom as having both a wave and particle character. The energy associated with the electrons is quantized or present in discrete energy packets.

There are 4 quantum numbers that bear directly on the position of electrons and their energy: The principle quantum number

varies with atomic number The azimuthal quantum which determines the orbital shape and angular momentum The magneto quantum number describes orientation of an orbital The spin quantum number describes electron spin

The following rules apply to orbitals

Rule

: Orbitals are designated s, p, d, and f and adhere to the following: s = spherical, 2 electrons p = sausage shape extending along x, y, and z axis, 6 electrons d = 5 degenerate orbitals along and between axes, 10 electrons f = (not a concern)

Rule

: At most, two electrons may occupy an orbital (or suborbital) and they must be of opposite spin

Rule

: s orbitals are spherical, with energy that varies only with distance from the nucleus. At most 2 electrons may occupy an S orbital.

Rule

: p orbitals extended along the major X, Y and Z axis designated

y and

z. Each holds 2 electrons, or 6 electrons to occupy the P orbital. The energy varies with both distance and direction

Rule

: d orbitals cover all space both along and between the axes. Their configuration is that of 5 degenerate (equal energy) and hold at most 10 electrons

The following rules apply to quantum states or atoms and orbitals

Rule

: Quantum states vary with atomic number, i.e., number of electrons

Rule

: Atoms with a principle quantum number

= 1 have only a 1

orbital. Examples are hydrogen and helium.

Rule

: Atoms with

= 2 have

and

orbitals

Rule

: Atoms with

= 3 have

, and

orbitals

Rule

: 4

orbitals are at a lower energy level than 3

and fill before 3

d Rule

: Atoms with 4

and 3

orbitals when ionizing lose 4

first

Two Major Rules in Chemical Physics that impinge on the behavior of minerals Hund’s rule: The lowest energy state of an atom is achieved when there is maximum utilization of the surrounding space by the occupying electrons. Pairing of electrons in an orbital is recognized as a higher energy state than single electrons of the same spin state occupying the orbitals. This does not apply to s orbitals.

Pauli exclusion principle: No two electrons in an atomic orbital may share the same set of quantum numbers. This rule led to the realization that electrons in the same orbital must be of opposite spins.

2pz

2py 2px 2s 1s 2s 3s 2px 2py 2pz

Quant No.

n=1 1s Configuration.

(K shell) n=2 n=3 2s, 2p (L shell) 3s, 3p, 3d (M shell)

Shapes are the same, but differ in orientation 2

orbitals. At the second quantum level orientation also becomes a factor in deciding orbital energy. Because there are 3 orientations existing simultaneously, a p orbital can hold a maximum of 6 electrons, 2 of opposite spin in each

Iron At. No. = 26 At. Wt.= 55.85

No. of occupying electrons

= 2

= 6

= 10

= 14

1

2

3

4

3

6 Principal Quantum Number (

=1, 2, 3)

Argon [Ar]4

3

6 Subshell (s,p,d,f)

Element (At. No.) Ground-state configuration Abbreviated form

Sodium (11) Magnesium (12) Aluminum (13) Silicon(14) 1

2 2

6 3

1 1

2 2

6 3

2 1

2 2

6 3

2 3

1 1

2 2

6 3

2 3

2 Phosphorus (15) Sulfur (16) Chlorine (17) Argon (18) 1

2 2

6 3

2 3

3 1

2 2

6 3

2 3

4 1

2 2

6 3

2 3

5 1

2 2

6 3

2 3

6 [Ne]3

1 [Ne]3

2 [Ne]3

2 3

1 [Ne]3

2 3

2 [Ne]3

2 3

3 [Ne]3

2 3

4 [Ne]3

2 3

5 [Ne]3

2 3

Caution:

The 4

orbital is actually at a lower energy level than the 3

. As a result 4

orbitals will fill before 3

. But, when ionized, electrons will be lost from the 4

before the 3

Class Exercise

Atomic numbers of

Potassium

and

Calcium

are 19 and 20, respectively. Outer electrons are in the M shell (

= 3). Determine the electronic configurations of potassium and calcium and determine their most likely ionized form

Solution:

When n = 3, the atom must contain s, p, and d subshells and 3 energy states. But, recall that the 4s subshell with 2 electrons is of a lower energy state than the 3d subshell and will fill first Z = 19

K = 1

2

3

4

Z = 20

Ca = 1

2

3

4

The most stable form occurs when both metals lose their 4s electrons. Thus:

K

and Ca

[Ar]4

and [Ar]4

Macrominerals Microminerals First transition series 3d 4d 5d

Elements in the First Transition Series

Sc Ti V Cr Mn Fe 3d 1 3d 2 3d 3 3d 5 3d 5 4s 2 4s 2 4s 2 4s 1 4s 2 3d 4s 6 2 Co Ni 3d 7 4s 2 Cu Zn 3d 8 3d 10 3d 10 4s 2 4s 1 4s 2 +3 +4 +3 +3 +2 +2 +4 +3 +3 +5 +4 +4 +1 +1 +1 +2 +2 +2 +3 +3 +2 Bio Essential Metals

1. Why do Ca 2+ , Mg 2+ , and Zn 2+ exist only as +2 ions? Li + , Na + and K + as +1 ions?

Ca = [Ar] 4s

Mg = [Ne] 3s

Zn = [Ar] 3d

4s

Li = [He] 2s

Na = [Ne] 3s

K = [Ar]4s

Octahedral Square planar Tetrahedral

element No. 26 1s

2s

2p

3s

3p

4s

3d

6 Z Z Z X

d xy

Y X

d X 2 -Y 2

Z X Y

d yz

X Y Z X Y

d Z 2 d xz

Ion Cu + Zn 2+ Cd 2+ Hg 2+ Coord Orb No.

10 4

10 4 2 sp sp sp 3 sp 3 3 Cu 2+ Ag 2+

9 4 4 dsp dsp 2 2

Metal Ion Antagonism

tetrahedral tetrahedral tetrahedral linear square planar square planar Prediction: Fe 2+

6 6 d 2 sp 3 octahedral Zn 2+ will interfere with Cu + Cd 2+ will interfere with Cu + and Zn 2+ Hg 2+ interference will be minimal Ag 2+ will interfere with Cu 2+ but not Zn 2+