-pure ionic compounds are formed by complete transfer of electrons from metallic atoms to non-metallic atoms. the electron cloud of an ion is solely confined to itself and not shared with its neighbouring ions. Actually pure ionic compounds are hard to find as the elctron cloud of an ion is easily distorted by the neighbouring oppositely charged ions and hence bonding has some covalent character.
(1) In pure ionic compound, the electron cloud is held tightly by the nucleus in an ion. No sharing of electrons is allowed between ions.
(2) In reality, if two ions of opposite charges are placed together, the electron cloud of an ion will be distorted by the other, thus inducing some sharing of electrons (covalency).
but for #1, the electron cloud is held tightly by the nucleus in an ion <-- how can one lose electrons and transfer to the other ?
Besides,
2. Explain the following characteristic properties of metals in terms of their structure and bonding
I) high density
II) high melting point
3. there are several crystalline forms of metallic iron. Below 906C, iron exists in the (alpha) form which has a boday centred cubic structure. Above this temperature, iron changes to (gama) form, which as a face centred cubic structure. Describe, with explanation, the change in colume that would occur when the temperature of a piece of iron is increased from room temperature to 100 C
4. which compound, AgCl, AgBr or AgI has the highest covalent character in the ionic bond. explain.
5. compare and explain the polar covalent bond of HCl, HBr, HI.
-----------
In part 1, it does not imply the full electron transfer. Just imagine a Li+ and a Cl- ions put nearby. The electron of cloud of Li+ is strongly held by the nucleus. So does the Cl nucleus to its electron cloud. However, the charge of Li+ will attract (though not strong enough to pull out electrons) the electron cloud of Cl-. This distortion of the theoretically spherical electron cloud (for the pure ionic model) accounts for the appearance of the covalent character of the so-called ionic bond.
2. Note that metal nuclei are held together by "electron sea", the free electrons that move around in metals. Therefore, metallic bonds are non-directional, which allows for the close packing of metal atoms. This results in the generally high density of metals.
Metallic bond is electrostatic in nature. Due to the efficient packing of atoms in metals, they usually have a high melting point (although this is not always the case; check the physical properties of group I and II elements).
3. Do you mean from RT to 1000C?
Alpha form of Fe is less packed (68%) compared to the gamma form of Fe (74%). Therefore, we would expect that the volume may reduce a bit during the phase transformation.
4. AgI should have the highest covalent character. It is due to the polarizability of the electron cloud of the halide ions whose order is:
F- < Cl- < Br- < I-
5. The bond length: HCl < HBr < HI
The dipole moment: HCl > HBr > HI
The anomaly arises from the increasing polarizability from Cl to I which results in the higher covalent character of the H-X bond down the group. The enhanced covalent character reduces the partial charges on H and X, thus decreasing the bond dipole moment.
----------------
1. X possibly belongs to group 1. Since X is polarizing (large charge density), the resulting X+Br- bond possesses a certain degree of covalent character which weakens the electrostatic attraction a bit. Therefore, the theoretical lattice enthalpy, which assumes a pure ionic interaction between X and Br, is larger than the experimental value.
2. It depends on whether the ionic character or the covalent character dominates.
3. Mg has two valence electrons while Na has only one. Therefore, the metallic bond between Mg atoms in solid is stronger than that of Na
4. Similar reason as 3. The melting point of group 1 elements decreases down the group because the atomic size gets larger which reduces the attraction between the nucleus and valence electron of an atom.
5. Metals are shiny since they reflect all lights. In more details, the free electrons of metals absorb light of all wavelengths and re-emit them. It means that all light is reflected by metals. Therefore, they have a lustre.
6. It means that the more negative side of a molecule will point toward the positive side of another molecule.
It depends on the liquid, but usually not since the induced dipole-ion interaction may not be strong enough the cause an observable attraction.
7. C-H bond should be polar since there is a difference in the electronegativity of C and H, but the bond dipole may not be large.
The dipole moment decreases from HCl to HBr to HI; therefore, we expect that the attraction between HCl molecules should be stronger than that between HBr molecules, and so do the HI molecules. Note that the stronger the attraction between molecules, the higher the boiling point.
However, this is not the case; experimentally, it is observed that the boiling point increases from HCl to HBr to HI. Obviously, there are other factoring coming into play to change the trend of the boiling point. The factor is the van der Waals force, which increases with the size and mass of the molecule. Since HI is much larger and heavier than HCl, the van der Waals force between HI molecules is so stronge that it compensates for the weak dipole-dipole attraction, resulting in an overall stronger attraction and thus the higher boiling point of HI compared to HBr and HCl.(new concept)
Sunday, August 30, 2009
the s-block elements
Chapter 9 The s-Block Elements
s-Block elements refer to elements having their outermost electrons occupying s-orbital. So, s-block elements also refer to group 1 and group 2 elements.
m.p. density crystal structure
Down a group : Li 181 0.53 b.c.c Be 1277 1.85 h.c.p
1. Increasing metallic character Na 98 0.97 b.c.c Mg 650 1.74 h.c.p
2. Weakening metallic bond K 64 0.86 b.c.c Ca 838 1.55 f.c.c
3. Increasing atomic size Rb 39 1.53 b.c.c Sr 768 2.6 f.c.c
4. Decreasing m.p. Cs 29 1.90 b.c.c Ba 714 3.5 b.c.c
9.1 Characteristic properties of the s-block elements
s-block elements are mainly metallic elements.
They have low electronegativities.
They form basic oxides and hydroxides.
They have fixed oxidation state in their compounds.
They do not have low lying d-orbitals to form complex.
Some of them have specific flame colour.
Li is anomalous to group 1 metal. Its chemistry resembles Mg rather than
a group 1 metal. This is part of the diagonal relationship.
Exercise 1. Why do s-block elements have low electronegativities?
2. Why do s-block elements have fixed oxidation states?
9.1.1 Metallic character
Group 1 and group 2 elements are metals.
- group 1 metals are named as "alkali metals".
- group 2 metals are named as "alkaline earth metals".
They are silvery and shiny but tarnish in air rapidly.
- freshly cut surface will expose metal to air and the metal
is oxidized.
They are low melting and can be cut by a knife.
- because each atom can only have 1 to 2 valence electrons to
form metallic bond. Of course, group 2 metals have
stronger metallic bond than group 1 metals.
Group 1 metals crystallize in body-centred cubic lattice.
- this explains the low density of group 1 metals
9.1.2 Electronegativity
The electronegativity of element
- decreases down a group and
- increases across a period
Explanation :
9.1.3 Formation of basic oxides and hydroxides
s-block elements have very low electronegativities and lose their
outermost electrons readily. Thus, they form ionic oxides and
ionic hydroxides which are necessarily basic in nature.
9.1.4 Standard electrode potential Eo /V Mn+(aq) + ne- → M(s)
Li Na K Rb Cs
-3.05 -2.71 -2.93 -2.99 -3.20
The abnormal value of Li is due to the exceptional small size of Li+
which is strongly hydrated to give the most negative Eo value.
Be Mg Ca Sr Ba
-1.85 -2.38 -2.87 -2.89 -2.90
9.1.5 Flame colour of compounds
Flame colour originates from the emission of visible light when an
excited electron de-excites(an electron returns to a lower energy level
from a higher energy level).
Before carrying out a flame test, the platinum wire(or silica rod) should
be washed with concentrated HCl by dipping the platinum wire into the
concentrated HCl and heat it in a non-luminous flame. As chlorides
are generally more volatile, the contaminant on the platinum wire
would be carried away. The cleaning process is completed when no
specific flame colour can be seen.
Flame colour of cations:
K+ lilac, viewed through a cobalt glass
Na+ golden yellow , persistent
Li+ crimson red, deep red
Ca2+ brick red
Sr2+ blood red
Ba2+ apple green
9.2 Variation in properties of the s-block elements and their compounds
9.2.1 Variation in atomic radii
Atomic radii increase down a group as the number of electron shell
is increasing and the outermost shell electrons are screened by core
electrons and experience a lower nuclear attraction.
Atomic radii decrease across a period because the addition electron
is put into the same quantum shell and the increase in nuclear charge
would attract the additional electron stronger to give a smaller atomic size.
atomic radius / nm
period number
9.2.2 Ionization Enthalpies
I.E. decreases down a group as the number of electron shell
is increasing and the outermost shell electrons are screened by core
electrons and experience a lower nuclear attraction.
First I.E. increases across a period because the addition electron
is put into the same quantum shell and the increase in nuclear charge
would attract the additional electron stronger to give a smaller atomic size.
The 2nd I.E. is always much greater than the 1st I.E. as it is more
difficult to remove a negatively charged electron from a positive ion.
9.2.3 Melting point
Plot the melting points of s-block elements in the following diagram.
m.p.(℃)
Period number
The melting point decreases down a group as the metallic bond is
weaker as the atomic size is larger.
The melting point of group 2 element is much larger than its group 1
counterpart as each group 2 atom has two electrons for bonding but
group 1 atom has one. The stronger metallic bond in group 2 metal
gives rise to higher melting point.
9.2.4 Hydration Enthalpies
The reaction of hydration is Mn+(g) + aq → Mn+(aq).
For an ion, hydration enthalpy is always negative.
The higher the charge density of an ion, the higher(more negative) the
hydration enthalpy. This is because the high electric potential holds
the water(polar solvent) molecules stronger.
hydration enthalpy kJ mol-1
-500 -2000
-200 -1500
Li+ Be2+
9.2.5 Reactions with oxygen
s-block elements reacts with oxygen gas readily.
K, Rb, Cs form superoxide, peroxide and normal oxide
Na, and Ba form peroxide and normal oxide
Others form normal oxides only.
K(s) + O2(g) → K2O(s) a normal oxide
K(s) + O2(g) → K2O2(s) a peroxide
K(s) + O2(g) → KO2(s) a superoxide
2-
potassium oxide K2O K+ O K+
2-
potassium peroxide K2O2 K+ O O K+
-
potassium superoxide KO2 K+ O O
-
potassium hydroxide KOH K+ O H
Soluble ionic oxide gives OH- and makes the solution alkaline.
2KO2(s) + 2H2O(l) → 2KOH(aq) + H2O2(aq) + 3O2(g)
2K2O2(s) + 2H2O(l) → 4KOH(aq) + O2(g)
K2O(s) + H2O(l) → KOH(aq)
9.2.6 Reactions with water
Sodium melts and moves rapidly on the surface of water. Hydrogen gas
is evolved and a strong alkaline solution is produced.
___ K(s) + ____H2O(l) →
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
____Ca(s) + ____H2O(l) →
____Mg(s) + ____H2O(g) →
____Al(s) + ____H2O(l) →
rocksil soaked
in water
heat heat magnesium
Exercise : Describe what you can see when a small piece of
(a) potassium
(b) magnesium
is dropped into water.
9.2.7 Reactions with hydrogen
Except Be , all s-block elements react with hydrogen.
Magnesium reacts under high pressure.
2K(s) + H2(g) 2KH(s) potassium hydride
Ca(s) + H2(g) CaH2(s)
The metal hydrides are ionic hydrides.
The hydride ion H- reacts with water vigorously.
KH(s) + H2O(l) → KOH(aq) + H2(g)
H-(s) + H2O(l) → OH-(aq) + H2(g)
9.2.8 Reactions with chlorine
All s-block elements react with chlorine.
2K(s) + Cl2(g) → 2KCl(s) potassium chloride
Ca(s) + Cl2(g) → CaCl2(s)
BeCl2(s) is covalent and has a linear polymeric structure.
9.2.9 Reactions of oxides, hydrides and chlorides with
water, acids and alkalis
metal oxide metal hydride metal chloride
water BeO(s) MgO(s), CaO(s)are insolubleOthers :K2O+H2O→2KOH(aq) CaH2+2H2O→Ca(OH)2+2H2 BeCl2+2H2O→Be(OH)2+2HClOthers no reaction
diluteacid NeutralizationMgO+2HCl→MgCl2+H2O Explosive BeCl2 reacts with waterOthers : no reaction
dilutealkali BeO+2NaOH+H2O→Na2Be(OH)4Others : no reaction Reacts with water in thealkaline solution. BeCl2+2NaOH→Be(OH)2+2NaClOthers : no reaction
9.2.10 Relative thermal stability of the carbonates and hydroxides
Stability of carbonates
carbonate
heat
lime water
Group 1 carbonates except Li2CO3 are stable to heat.
Group 2 carbonates decompose on heating.
Decomposition temperature:
BeCO3 MgCO3 CaCO3 SrCO3 BaCO3
~100℃ 540 900 1290 1360
Group 2 carbonates are becoming less stable to heat on going down the group.
BaCO3 is quite stable to heat.
MgCO3 is most unstable to heat because its cation is far more smaller than the
anion. The small cation has a high charge density and high polarizing power
and polarizes the large anion. The anion is distorted and can be decomposed
easily.
Stability of hydroxides
Li and group 2 metal hydroxides are decomposable at high temperature.
The ease of decomposition decreases down group 2. The trend again is
governed by the polarizing power of cation.
Mg(OH)2(s) MgO(s) + H2O(g)
9.2.11 Relative solubility of the sulphates(VI) and hydroxides
Sulphates
Group 1 sulphates are soluble.
Group 2 sulphates :
Solubility of sulphates : MgSO4 > CaSO4 > SrSO4 > BaSO4
soluble insoluble insoluble insoluble
Reason : Sulphate is a large anion. The size of the cations down the
group are progressively increasing. The decrease in
hydration enthalpy of the cation outweighs the decrease in
lattice enthalpy. So, sulphates become less soluble down the group.
MgSO4(s) + aq → MgSO4(aq) enthalpy of solution of MgSO4
-lattice energy hydration enthalpy of Mg2+(g)
+ hydration enthalpy of SO42-(g)
Mg2+(g) + SO42-(g)
Hydroxides
Group 1 hydroxides are strong alkalis.
Group 2 hydroxides :
Solubility of hydroxides :
Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2
insoluble insoluble soluble soluble
Reason : __________ is a small anion. The size of the cations down
the group are progressively increasing. The decrease in
____________ enthalpy of the cation outweighs the decrease
in ___________ enthalpy. So, hydroxides become ________
soluble down the group.
Energy cycle :
9.2.12 Uses of s-block elements
1. Na2CO3 is useful in making soda glass.
CaCO3(s) + SiO2(s) CaSiO3(s) + CO2(g)
Na2CO3(s) + SiO2(s) Na2SiO3(s) + CO2(g)
sodium silicate
Soda glass is a mixture of calcium silicate and sodium silicate.
2. NaHCO3 as baking powder
2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g)
The carbon dioxide evolved will rise the cake.
3. NaOH/KOH in making soap
Fat molecule can be hydrolysed(saponification) to form soap.
e.g. glyceryl stearate + alkali → glycerol + sodium stearate(soap)
4. Mg(OH)2 as anti-acid to neutralize excess acid in stomach
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
5. Slaked lime to neutralize acid in industrial effluent
Ca(OH)2(s) + 2HCl(aq) → CaCl2(aq) + H2O(l)
6. Strontium in firework
The burning of strontium compounds gives ___________ _________ colour glare.
s-Block elements refer to elements having their outermost electrons occupying s-orbital. So, s-block elements also refer to group 1 and group 2 elements.
m.p. density crystal structure
Down a group : Li 181 0.53 b.c.c Be 1277 1.85 h.c.p
1. Increasing metallic character Na 98 0.97 b.c.c Mg 650 1.74 h.c.p
2. Weakening metallic bond K 64 0.86 b.c.c Ca 838 1.55 f.c.c
3. Increasing atomic size Rb 39 1.53 b.c.c Sr 768 2.6 f.c.c
4. Decreasing m.p. Cs 29 1.90 b.c.c Ba 714 3.5 b.c.c
9.1 Characteristic properties of the s-block elements
s-block elements are mainly metallic elements.
They have low electronegativities.
They form basic oxides and hydroxides.
They have fixed oxidation state in their compounds.
They do not have low lying d-orbitals to form complex.
Some of them have specific flame colour.
Li is anomalous to group 1 metal. Its chemistry resembles Mg rather than
a group 1 metal. This is part of the diagonal relationship.
Exercise 1. Why do s-block elements have low electronegativities?
2. Why do s-block elements have fixed oxidation states?
9.1.1 Metallic character
Group 1 and group 2 elements are metals.
- group 1 metals are named as "alkali metals".
- group 2 metals are named as "alkaline earth metals".
They are silvery and shiny but tarnish in air rapidly.
- freshly cut surface will expose metal to air and the metal
is oxidized.
They are low melting and can be cut by a knife.
- because each atom can only have 1 to 2 valence electrons to
form metallic bond. Of course, group 2 metals have
stronger metallic bond than group 1 metals.
Group 1 metals crystallize in body-centred cubic lattice.
- this explains the low density of group 1 metals
9.1.2 Electronegativity
The electronegativity of element
- decreases down a group and
- increases across a period
Explanation :
9.1.3 Formation of basic oxides and hydroxides
s-block elements have very low electronegativities and lose their
outermost electrons readily. Thus, they form ionic oxides and
ionic hydroxides which are necessarily basic in nature.
9.1.4 Standard electrode potential Eo /V Mn+(aq) + ne- → M(s)
Li Na K Rb Cs
-3.05 -2.71 -2.93 -2.99 -3.20
The abnormal value of Li is due to the exceptional small size of Li+
which is strongly hydrated to give the most negative Eo value.
Be Mg Ca Sr Ba
-1.85 -2.38 -2.87 -2.89 -2.90
9.1.5 Flame colour of compounds
Flame colour originates from the emission of visible light when an
excited electron de-excites(an electron returns to a lower energy level
from a higher energy level).
Before carrying out a flame test, the platinum wire(or silica rod) should
be washed with concentrated HCl by dipping the platinum wire into the
concentrated HCl and heat it in a non-luminous flame. As chlorides
are generally more volatile, the contaminant on the platinum wire
would be carried away. The cleaning process is completed when no
specific flame colour can be seen.
Flame colour of cations:
K+ lilac, viewed through a cobalt glass
Na+ golden yellow , persistent
Li+ crimson red, deep red
Ca2+ brick red
Sr2+ blood red
Ba2+ apple green
9.2 Variation in properties of the s-block elements and their compounds
9.2.1 Variation in atomic radii
Atomic radii increase down a group as the number of electron shell
is increasing and the outermost shell electrons are screened by core
electrons and experience a lower nuclear attraction.
Atomic radii decrease across a period because the addition electron
is put into the same quantum shell and the increase in nuclear charge
would attract the additional electron stronger to give a smaller atomic size.
atomic radius / nm
period number
9.2.2 Ionization Enthalpies
I.E. decreases down a group as the number of electron shell
is increasing and the outermost shell electrons are screened by core
electrons and experience a lower nuclear attraction.
First I.E. increases across a period because the addition electron
is put into the same quantum shell and the increase in nuclear charge
would attract the additional electron stronger to give a smaller atomic size.
The 2nd I.E. is always much greater than the 1st I.E. as it is more
difficult to remove a negatively charged electron from a positive ion.
9.2.3 Melting point
Plot the melting points of s-block elements in the following diagram.
m.p.(℃)
Period number
The melting point decreases down a group as the metallic bond is
weaker as the atomic size is larger.
The melting point of group 2 element is much larger than its group 1
counterpart as each group 2 atom has two electrons for bonding but
group 1 atom has one. The stronger metallic bond in group 2 metal
gives rise to higher melting point.
9.2.4 Hydration Enthalpies
The reaction of hydration is Mn+(g) + aq → Mn+(aq).
For an ion, hydration enthalpy is always negative.
The higher the charge density of an ion, the higher(more negative) the
hydration enthalpy. This is because the high electric potential holds
the water(polar solvent) molecules stronger.
hydration enthalpy kJ mol-1
-500 -2000
-200 -1500
Li+ Be2+
9.2.5 Reactions with oxygen
s-block elements reacts with oxygen gas readily.
K, Rb, Cs form superoxide, peroxide and normal oxide
Na, and Ba form peroxide and normal oxide
Others form normal oxides only.
K(s) + O2(g) → K2O(s) a normal oxide
K(s) + O2(g) → K2O2(s) a peroxide
K(s) + O2(g) → KO2(s) a superoxide
2-
potassium oxide K2O K+ O K+
2-
potassium peroxide K2O2 K+ O O K+
-
potassium superoxide KO2 K+ O O
-
potassium hydroxide KOH K+ O H
Soluble ionic oxide gives OH- and makes the solution alkaline.
2KO2(s) + 2H2O(l) → 2KOH(aq) + H2O2(aq) + 3O2(g)
2K2O2(s) + 2H2O(l) → 4KOH(aq) + O2(g)
K2O(s) + H2O(l) → KOH(aq)
9.2.6 Reactions with water
Sodium melts and moves rapidly on the surface of water. Hydrogen gas
is evolved and a strong alkaline solution is produced.
___ K(s) + ____H2O(l) →
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
____Ca(s) + ____H2O(l) →
____Mg(s) + ____H2O(g) →
____Al(s) + ____H2O(l) →
rocksil soaked
in water
heat heat magnesium
Exercise : Describe what you can see when a small piece of
(a) potassium
(b) magnesium
is dropped into water.
9.2.7 Reactions with hydrogen
Except Be , all s-block elements react with hydrogen.
Magnesium reacts under high pressure.
2K(s) + H2(g) 2KH(s) potassium hydride
Ca(s) + H2(g) CaH2(s)
The metal hydrides are ionic hydrides.
The hydride ion H- reacts with water vigorously.
KH(s) + H2O(l) → KOH(aq) + H2(g)
H-(s) + H2O(l) → OH-(aq) + H2(g)
9.2.8 Reactions with chlorine
All s-block elements react with chlorine.
2K(s) + Cl2(g) → 2KCl(s) potassium chloride
Ca(s) + Cl2(g) → CaCl2(s)
BeCl2(s) is covalent and has a linear polymeric structure.
9.2.9 Reactions of oxides, hydrides and chlorides with
water, acids and alkalis
metal oxide metal hydride metal chloride
water BeO(s) MgO(s), CaO(s)are insolubleOthers :K2O+H2O→2KOH(aq) CaH2+2H2O→Ca(OH)2+2H2 BeCl2+2H2O→Be(OH)2+2HClOthers no reaction
diluteacid NeutralizationMgO+2HCl→MgCl2+H2O Explosive BeCl2 reacts with waterOthers : no reaction
dilutealkali BeO+2NaOH+H2O→Na2Be(OH)4Others : no reaction Reacts with water in thealkaline solution. BeCl2+2NaOH→Be(OH)2+2NaClOthers : no reaction
9.2.10 Relative thermal stability of the carbonates and hydroxides
Stability of carbonates
carbonate
heat
lime water
Group 1 carbonates except Li2CO3 are stable to heat.
Group 2 carbonates decompose on heating.
Decomposition temperature:
BeCO3 MgCO3 CaCO3 SrCO3 BaCO3
~100℃ 540 900 1290 1360
Group 2 carbonates are becoming less stable to heat on going down the group.
BaCO3 is quite stable to heat.
MgCO3 is most unstable to heat because its cation is far more smaller than the
anion. The small cation has a high charge density and high polarizing power
and polarizes the large anion. The anion is distorted and can be decomposed
easily.
Stability of hydroxides
Li and group 2 metal hydroxides are decomposable at high temperature.
The ease of decomposition decreases down group 2. The trend again is
governed by the polarizing power of cation.
Mg(OH)2(s) MgO(s) + H2O(g)
9.2.11 Relative solubility of the sulphates(VI) and hydroxides
Sulphates
Group 1 sulphates are soluble.
Group 2 sulphates :
Solubility of sulphates : MgSO4 > CaSO4 > SrSO4 > BaSO4
soluble insoluble insoluble insoluble
Reason : Sulphate is a large anion. The size of the cations down the
group are progressively increasing. The decrease in
hydration enthalpy of the cation outweighs the decrease in
lattice enthalpy. So, sulphates become less soluble down the group.
MgSO4(s) + aq → MgSO4(aq) enthalpy of solution of MgSO4
-lattice energy hydration enthalpy of Mg2+(g)
+ hydration enthalpy of SO42-(g)
Mg2+(g) + SO42-(g)
Hydroxides
Group 1 hydroxides are strong alkalis.
Group 2 hydroxides :
Solubility of hydroxides :
Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2
insoluble insoluble soluble soluble
Reason : __________ is a small anion. The size of the cations down
the group are progressively increasing. The decrease in
____________ enthalpy of the cation outweighs the decrease
in ___________ enthalpy. So, hydroxides become ________
soluble down the group.
Energy cycle :
9.2.12 Uses of s-block elements
1. Na2CO3 is useful in making soda glass.
CaCO3(s) + SiO2(s) CaSiO3(s) + CO2(g)
Na2CO3(s) + SiO2(s) Na2SiO3(s) + CO2(g)
sodium silicate
Soda glass is a mixture of calcium silicate and sodium silicate.
2. NaHCO3 as baking powder
2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g)
The carbon dioxide evolved will rise the cake.
3. NaOH/KOH in making soap
Fat molecule can be hydrolysed(saponification) to form soap.
e.g. glyceryl stearate + alkali → glycerol + sodium stearate(soap)
4. Mg(OH)2 as anti-acid to neutralize excess acid in stomach
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
5. Slaked lime to neutralize acid in industrial effluent
Ca(OH)2(s) + 2HCl(aq) → CaCl2(aq) + H2O(l)
6. Strontium in firework
The burning of strontium compounds gives ___________ _________ colour glare.
Thursday, March 26, 2009
atomic structure
Orbitals, s, p, d, f and average distance from the nucleus
Orbitals are the regions of space occupied by electrons in atoms. The shapes and volumes of space are different for each type of orbital. The electrons are confined to specific shaped spaces around the nucleus. The electrons are not on the surface of these shapes. The electrons are moving inside the volume marked out by the shape.
The shape for all "s" orbitals is spherical with the center of the sphere at the nucleus. The size of the sphere increases for increasing quantum numbers. The 1s has a smaller diameter than the 2s and so forth.
1s <>
The p orbitals have a double lobe shape with the point of contact of the lobes at the center of the nucleus. The three different p orbitals are each aligned along one of the three coordinate axes x, y, or z.
The principle quantum number "n" is very useful. It tells the relative energy for an electron in an atom and it also indicates the average distance between the nucleus and an electron. The higher the value for "n" the greater the distance between the nucleus and the electron. An electron in the first level, n = 1, is closer to the nucleus on average than an electron in the n = 4 shell. Exercise: What do the subscripts x,y an z indicate about the "p" orbitals in a shell? Click here for answer
Number of orbitals in a shell
The number of each type of orbital has a mathematical origin. You need to know only the results of the math.
orbital type s p d f
Number of orbitals in a shell
The number of each type of orbital has a mathematical origin. You need to know only the results of the math.
orbital type s p d f
maximum number of orbitals in a shell 1 3 5 7
number of electrons that can fit into orbital type 2 6 10 14
number of electrons that can fit into orbital type 2 6 10 14
Notice that the number of orbitals increases by "2" in going from one type to another. The capacity for electrons increases by "4" at the same time.
Exercise:
How many "g" orbitals would you predict? Answer: There should be 9 and they could hold 18 electrons.
Number of electrons in an orbital and electron spin
Every orbital can only hold two electrons. This is linked to the magnetic character of electrons. The electrons act as though they are spinning charged particles. There are two spin directions that match the two magnetic fields created by electrons. There is a spin up field and a spin down field. Commonly the two magnetic fields and spin types are indicated by arrows. An up arrow for spin up and a down arrow for spin down. The pairing of the magnetic fields of two electrons cancels out the fields and sets the limit on the population of an orbital at two.
Maximum number of electrons in any orbital = 2
Atoms with all the electron spins paired have no magnetic character. An atom with unpaired electrons shows magnetic properties. This is the origin of the magnetic behavior of iron and nickel for example. This magnetic property gives us a way to know how the electrons are arranged in atoms. Particles with unpaired electrons are called as "free radicals". Free radicals have been linked to the aging process, mutations in DNA, and cancer.
Exercise: How many electrons can fit into all of the 3p subshell orbitals? Click here for answer.
Exercise:
How many "g" orbitals would you predict? Answer: There should be 9 and they could hold 18 electrons.
Number of electrons in an orbital and electron spin
Every orbital can only hold two electrons. This is linked to the magnetic character of electrons. The electrons act as though they are spinning charged particles. There are two spin directions that match the two magnetic fields created by electrons. There is a spin up field and a spin down field. Commonly the two magnetic fields and spin types are indicated by arrows. An up arrow for spin up and a down arrow for spin down. The pairing of the magnetic fields of two electrons cancels out the fields and sets the limit on the population of an orbital at two.
Maximum number of electrons in any orbital = 2
Atoms with all the electron spins paired have no magnetic character. An atom with unpaired electrons shows magnetic properties. This is the origin of the magnetic behavior of iron and nickel for example. This magnetic property gives us a way to know how the electrons are arranged in atoms. Particles with unpaired electrons are called as "free radicals". Free radicals have been linked to the aging process, mutations in DNA, and cancer.
Exercise: How many electrons can fit into all of the 3p subshell orbitals? Click here for answer.
atomic structure
Total enegies of an atom: potential energy and kinetic energy.
An unusual feature of electrons and atoms is that the energy is "quantized" or restricted to definite values. This is not what we see on a daily basis with large objects. Big things appear to have continuous energy states, with no restrictions or jumps. The emission spectra of atoms provides the proof of the quantum theory as it applies to atoms and subatomic particles. Every element can be made to emit light that is characteristic for that element. This is consistent with the idea that each element has a unique energy level pattern. The light emitted by an element's atoms matches the possible jumps betweeen energy levels for the electrons in the atoms. The total energy for a specific shell is restricted. The total can be made up of a high kinetic energy(KE) and a low potential energy(PE). Sometimes the KE is low and the potential energy is high. The total is what counts.
Energy total = kinetic energy + potential energy The basic idea is that the electrons in atoms can only have energy values governed by the principle quantum number "n". The principle quantum number can have only positive whole number values. This means no fractions and no negative numbers. The allowed values for "n" are 1, 2, 3, up to infinity. Electrons in atoms typically only use the first seven energy levels.
Number of electrons in a shell or level
The principle quantum number "n" is very useful. It provides a way to dtermine the maximum number of elecytrons that can fit into a shell. The number of electrons that can fit into a shell is given by the following formula.
Maximum number of electrons in level = 2n2
This means that the first shell can hold only two electrons. The second shell with n = 2 can hold only eight, 8, electrons.
Row in periodic table
Shell
2n2
Maximum number of electrons in shell
1
n=1
2(1)2 = 2
2
2
n=2
2(2)2= 8
8
3
n=3
2(3)2= 18
18
Distance between nucleus and an electron
The principle quantum number "n" is very useful. It tells the relative energy for an electron in an atom and it also indicates the average distance between the nucleus and an electron. The higher the value for "n" the greater the distance between the nucleus and the electron. An electron in the first level, n = 1, is closer to the nucleus on average than an electron in the n = 4 shell.
Shells, number of subshells, and order of filling
The shells or energy levels are also divided into sublevels or subshells. The labels for these subshells comes from the observations of light emitted by atoms. The labels for subshells are s for sharp, p for principle, d for diffuse, f for fine. The physical connections end with the f sublevel class. The relative energies for these sublevels are as follows.
s < p < d < f < g < etc.
The value for "n" equals the number of sublevels for a level. This means the first shell or level has only one sublevel. The second level with "n" equal to two has two sublevels. Check the table below to see what happens for big "n" values. Thankfully, the number of electrons (112) in the biggest atom known can be accommodated by using only the common s, p, d, and f sublevels. The subshells g, h, i etc. are never needed.
level or shell
number of subshells
type of subshells
n = 1
1
s
n = 2
2
s and p
n = 3
3
s, p, d
n = 4
4
s, p, d, f
n = 5
5
s, p, d, f, g
n = 6
6
s, p, d, f, g, h
n = 7
7
s, p, d, f, g, h, i
Electrons behave like they are small magnets. They act like they are spinning like a top. The direction of the magnetic field generated by the "spinning " electron is limited to two extremes. The "UP" and the "DOWN" conditions. The electron spin can be cancelled when two electrons of opposite spin are in the same orbital as happens in helium, He 1s2. The helium atom shows no magnetic properties. Atoms with an odd number of electrons like lithium, Li 1s2 2s1 are paramagnetic.
Atoms and ions with unpaired electrons are "paramagnetic". They exert a magnetic field. Atoms with many unpaired electrons exert a stronger magnetic field. These observable properties are one piece of evidence for our model of electron configurations. In metals like iron the unpaired electrons act in concert and form "domains" with an increased magnetic character. This is called ferromagnetism.
Free radicals are particles that have an unpaired electron. Their existence is detected by checking for magnetic fields. The radicals are reactive and are believed to play a role in aging and initiation of cancer. Antioxidants are molecules that react with free radicals and remove them. This keeps them from reacting with normal molecules in the body.
electron spin produces magnetism
An unusual feature of electrons and atoms is that the energy is "quantized" or restricted to definite values. This is not what we see on a daily basis with large objects. Big things appear to have continuous energy states, with no restrictions or jumps. The emission spectra of atoms provides the proof of the quantum theory as it applies to atoms and subatomic particles. Every element can be made to emit light that is characteristic for that element. This is consistent with the idea that each element has a unique energy level pattern. The light emitted by an element's atoms matches the possible jumps betweeen energy levels for the electrons in the atoms. The total energy for a specific shell is restricted. The total can be made up of a high kinetic energy(KE) and a low potential energy(PE). Sometimes the KE is low and the potential energy is high. The total is what counts.
Energy total = kinetic energy + potential energy The basic idea is that the electrons in atoms can only have energy values governed by the principle quantum number "n". The principle quantum number can have only positive whole number values. This means no fractions and no negative numbers. The allowed values for "n" are 1, 2, 3, up to infinity. Electrons in atoms typically only use the first seven energy levels.
Number of electrons in a shell or level
The principle quantum number "n" is very useful. It provides a way to dtermine the maximum number of elecytrons that can fit into a shell. The number of electrons that can fit into a shell is given by the following formula.
Maximum number of electrons in level = 2n2
This means that the first shell can hold only two electrons. The second shell with n = 2 can hold only eight, 8, electrons.
Row in periodic table
Shell
2n2
Maximum number of electrons in shell
1
n=1
2(1)2 = 2
2
2
n=2
2(2)2= 8
8
3
n=3
2(3)2= 18
18
Distance between nucleus and an electron
The principle quantum number "n" is very useful. It tells the relative energy for an electron in an atom and it also indicates the average distance between the nucleus and an electron. The higher the value for "n" the greater the distance between the nucleus and the electron. An electron in the first level, n = 1, is closer to the nucleus on average than an electron in the n = 4 shell.
Shells, number of subshells, and order of filling
The shells or energy levels are also divided into sublevels or subshells. The labels for these subshells comes from the observations of light emitted by atoms. The labels for subshells are s for sharp, p for principle, d for diffuse, f for fine. The physical connections end with the f sublevel class. The relative energies for these sublevels are as follows.
s < p < d < f < g < etc.
The value for "n" equals the number of sublevels for a level. This means the first shell or level has only one sublevel. The second level with "n" equal to two has two sublevels. Check the table below to see what happens for big "n" values. Thankfully, the number of electrons (112) in the biggest atom known can be accommodated by using only the common s, p, d, and f sublevels. The subshells g, h, i etc. are never needed.
level or shell
number of subshells
type of subshells
n = 1
1
s
n = 2
2
s and p
n = 3
3
s, p, d
n = 4
4
s, p, d, f
n = 5
5
s, p, d, f, g
n = 6
6
s, p, d, f, g, h
n = 7
7
s, p, d, f, g, h, i
Electrons behave like they are small magnets. They act like they are spinning like a top. The direction of the magnetic field generated by the "spinning " electron is limited to two extremes. The "UP" and the "DOWN" conditions. The electron spin can be cancelled when two electrons of opposite spin are in the same orbital as happens in helium, He 1s2. The helium atom shows no magnetic properties. Atoms with an odd number of electrons like lithium, Li 1s2 2s1 are paramagnetic.
Atoms and ions with unpaired electrons are "paramagnetic". They exert a magnetic field. Atoms with many unpaired electrons exert a stronger magnetic field. These observable properties are one piece of evidence for our model of electron configurations. In metals like iron the unpaired electrons act in concert and form "domains" with an increased magnetic character. This is called ferromagnetism.
Free radicals are particles that have an unpaired electron. Their existence is detected by checking for magnetic fields. The radicals are reactive and are believed to play a role in aging and initiation of cancer. Antioxidants are molecules that react with free radicals and remove them. This keeps them from reacting with normal molecules in the body.
electron spin produces magnetism
A beam of rays came out from the cathode and hit the anode is called cathode..
The beam was composed of negatively charged fast-moving particles.
•Most of the mass of the atom was carried by the electrons (>1000 e-)
• An atom was a positively charged sphere with low density
•Negatively charged electrons embedded in it like a ‘plum pudding’
The beam was composed of negatively charged fast-moving particles.
•Most of the mass of the atom was carried by the electrons (>1000 e-)
• An atom was a positively charged sphere with low density
•Negatively charged electrons embedded in it like a ‘plum pudding’
Tuesday, March 24, 2009
determine the conc.
During today's chem lesson, Miss Jee asked us to think about how to determine the concentration of the solution.
a. Students will create a standard curve, plotting the relationship between the concentration of a solution and the amount of light that it is able to absorb. [Content Standard Unifying Concepts- Change, constancy, and measurement] [9-12 Content Standard B- Properties of matter]
b. Students will be able to use the standard curve they have created to determine the concentration of an unknown solution. [9-12 Content Standard A- Formulate explanations using evidence]
Materials:
CBL System
TI Graphing Calculator [9-12 Content Standard E- Understandings about science and technology]
Vernier Colorimeter
Vernier adapter cable
TI-Graph Link
One cuvette
Five 20 x 150 mm test tubes
Tissues (preferrably lint free)
30 mL of 0.40 M NiSO4
5 mL of NiSO4 of unknown concentration
two 10 ml pipets (or graduated cylinders)
pipet pump or pipet bulb
distilled water
test tube rack
two 100mL beakers
stirring rod [Teaching Standard D- Make accessible science tools]
Procedure:
1. Using the table below, use the distilled water provided to dilute the NiSO4 and make 5 solutions with known concentrations.
Trial Number
0.40 M NiSO4
(mL)
H20
(mL)
Concentration
(M)
1
2
8
0.08
2
4
6
0.16
3
6
4
0.24
4
8
2
0.32
5
10
0
0.40
2. Calibrate the colorimeter. Prepare a blank by filling a cuvette ¾ full of distilled water. With the light source turned off, enter this absorbance value obtained as 0% transmittance. With the wavelength knob in the Red LED position (635 nm), enter the absorbance value obtained as 100% transmittance.
3. In this same manner, collect absorbance data for each of the five standard solutions. When the percent transmittance value for each solution is displayed , enter the molar concentration for that solution.
4. Using your calculator, construct a graph of absorbance vs. concentration. Then perform a linear regression on your data. [9-12 Content Standard A- Use mathematics to improve scientific communication] If the data you have obtained are consistent with Beer’s Law (a direct relationship between absorbance and concentration), the regression line should closely fit the five data points and should pass through (or near) the origin of the graph.
5. Obtain about 5 mL of the unknown solution of NiSO4 . Find the absorbance for the unknown solution. Then use your calculator to interpolate along the regression line on your Beer’s Law curve.
6. Use the TI Graph link cable and program to transfer the graph of absorbance vs. concentration (including the interpolated unknown concentration) to a laptop computer. Print a copy of the graph.
b. Students will be able to use the standard curve they have created to determine the concentration of an unknown solution. [9-12 Content Standard A- Formulate explanations using evidence]
Materials:
CBL System
TI Graphing Calculator [9-12 Content Standard E- Understandings about science and technology]
Vernier Colorimeter
Vernier adapter cable
TI-Graph Link
One cuvette
Five 20 x 150 mm test tubes
Tissues (preferrably lint free)
30 mL of 0.40 M NiSO4
5 mL of NiSO4 of unknown concentration
two 10 ml pipets (or graduated cylinders)
pipet pump or pipet bulb
distilled water
test tube rack
two 100mL beakers
stirring rod [Teaching Standard D- Make accessible science tools]
Procedure:
1. Using the table below, use the distilled water provided to dilute the NiSO4 and make 5 solutions with known concentrations.
Trial Number
0.40 M NiSO4
(mL)
H20
(mL)
Concentration
(M)
1
2
8
0.08
2
4
6
0.16
3
6
4
0.24
4
8
2
0.32
5
10
0
0.40
2. Calibrate the colorimeter. Prepare a blank by filling a cuvette ¾ full of distilled water. With the light source turned off, enter this absorbance value obtained as 0% transmittance. With the wavelength knob in the Red LED position (635 nm), enter the absorbance value obtained as 100% transmittance.
3. In this same manner, collect absorbance data for each of the five standard solutions. When the percent transmittance value for each solution is displayed , enter the molar concentration for that solution.
4. Using your calculator, construct a graph of absorbance vs. concentration. Then perform a linear regression on your data. [9-12 Content Standard A- Use mathematics to improve scientific communication] If the data you have obtained are consistent with Beer’s Law (a direct relationship between absorbance and concentration), the regression line should closely fit the five data points and should pass through (or near) the origin of the graph.
5. Obtain about 5 mL of the unknown solution of NiSO4 . Find the absorbance for the unknown solution. Then use your calculator to interpolate along the regression line on your Beer’s Law curve.
6. Use the TI Graph link cable and program to transfer the graph of absorbance vs. concentration (including the interpolated unknown concentration) to a laptop computer. Print a copy of the graph.
career
Mm.. Thank you for the comments and i would like the fully explain what this career is about after i interview a shipping broker:)
First of all, a shipping broker doesn't need to swim or even go anywhere near the sea. Being a broker, all i have to do is to use the computer and exchange information of the shipping industry, for example, which buyer is interested and which ship is available to carry the goods.. Then, shipping brokers can help different parties to cooperate. Basically, communication is very important so shipping brokers normally need to talk a lot on the phone and use msn a lot. In addition, networking for them is also very important. The other reason I like this job is that you get to travel to different places to meet the clients around the world and you can also have high-class because you need to accompany the clients. Because they need to work with clients world-wide, they don’t get to sleep a lot and that’s the part I don’t like about this job.
You might like to ask yourself the following questions:
1) What are some of my strengths?
Maybe is my communication ability because I can easily talk with strangers and I like to talk on the phone a lot.
2) What are some of my weaknesses?
I really need to sleep on time.
3) What are the opportunities available in the short term / long term?
I don’t quite understand the question..
4) How about threats or potential problems that you foresee in this job/industry?
I don’t think there is any because the world is more globalize and there is a great need of transportation of goods.
Subscribe to:
Posts (Atom)