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16.4.09
Recaps.
Metallic Bonding
1)-The electromagnetic interaction between delocalized electrons, called conduction electrons, and the metallic nuclei within metals.
-Electropositive elements.
Understood as the sharing of "free" electrons among a lattice of positively-charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals.
2)In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities.
3)Metallic bonding accounts for many physical properties of metals,such as:
-High Melting and Boiling point,(Very strong electrostatic forces of attraction between sea of electrons and the positive metal ions.)
-Strength,
-Malleability and Ductility,(Due to orderly packing of metal atoms.The layers of atom can slide over each other easily whenever a force is applied.)
-Thermal and Electrical conductivity,(They have mobile electrons to carry electricity and heat energy.)
-Opacity, and
-Luster.
4)Although the term metallic bond is often used in contrast to the term covalent bond, it is more preferable to use the term metallic bonding, because this type of bonding is collective in nature and a single "metallic bond" does not exist.
5)The atoms of a metal contribute their valence electrons to form a 'sea' of electrons.
6)The sea of electrons helps to cement the positive ions together to give metals a giant metallic structure.
Physical Properties
1)The physical properties of a substance are determined by its structure and bonding.
STRUCTURE=>PARTICLES=>BONDING BETWEEN PARTICLES=>M.P
Simple Covalent->Simple molecules->Weak Van der waals F.O.A->Low
Macromolecules->Atoms->Strong covalent bonds->High
Metallic->+Ions in "sea" of e- =>Strong metallic Bonds->High
Ionic->Ions->Strong Ionic Bonds->High
The nature of metallic bonding
-The combination of two phenomena gives rise to metallic bonding: delocalization of electrons and the availability of a far larger number of delocalized energy states than of delocalized electrons. The latter could be called electron deficiency.
Delocalization
In 2D
Delocalization -bonding that involves more than one pair of atoms held together by one pair of electrons- is most familiar from the example of benzene C6H6, where six electrons from six carbon atoms are engaged in joint aromatic bonding. However, there are other examples like the three-center two-electron bond, prevalent in boron chemistry. The principle can easily be extended over larger aromatic molecules like naphthalene, anthracene and if the process is taken to its extreme: graphite. The latter is an example of a system delocalized in two dimensions. Interestingly, there is an isoelectronic analog of benzene, B3N3H6 (borazine) for which the same argument holds. It has very similar properties to benzene[5] When extended infinitely hexagonal boron nitride BN is obtained with a structure identical to that of graphite, apart from the alternation between boron and nitrogen in each ring. This material is a semiconductor, exemplifying that delocalization is a necessary but not sufficient requirement for conductivity. Electrical conductivity does occur in graphite, because the π and π*-like bands overlap, making it a semimetal, with partly filled bands, fulfilling the other requirement for conductivity.
In 3D
Metal aromaticity in metal clusters is another example of delocalization, this time often in three-dimensional entities. Metals take the delocalization principle to its extreme and one could say that a crystal of a metal represents a single molecule over which all conduction electrons are delocalized in all three dimensions. This means that inside the metal one can generally not distinguish molecules so that the metallic bonding is neither intra- nor intermolecular. 'Nonmolecular' would perhaps be a better term. Metallic bonding is mostly non-polar, because even in alloys there is little difference among the electronegativities of the atoms participating in the bonding interaction (and in pure elemental metals, none at all). Thus metallic bonding is an extremely delocalized communal form of covalent bonding. In a sense metallic bonding is not a 'new' type of bonding at all therefore and it only describes the bonding as present in a chunk of condensed matter, be it crystalline solid, liquid or even glass. Metallic vapors by contrast are often atomic (Hg) or at times contain molecules like Na2 held together by a more conventional covalent bond. This is why it is not correct to speak of a single 'metallic bond'.
The delocalization is most pronounced for s- and p-electrons. For cesium it is so strong that the electrons are virtually free from the cesium atoms to form a gas only constrained by the surface of the metal. For cesium therefore the picture of Cs+-ions held together by a negatively charged electron gas is not too inaccurate [6]. For other elements the electrons are less free, in that they still experience the potential of the metal atoms, sometimes quite strongly. They require a more intricate quantum mechanical treatment (e.g. tight binding) in which the atoms are viewed as neutral much like the carbon atoms in benzene. For d- and especially f-electrons the delocalization is not strong at all and this explains why these electrons are able to continue behaving as unpaired electrons that retain their spin, adding interesting magnetic properties to these metals.
Electron deficiency and mobility
Metal atoms contain few electrons in their valence shells relative to their periods or energy levels. They are electron deficient elements and the communal sharing does not change that. There remain far more available energy states than there are shared electrons. Both requirements for conductivity are therefore fulfilled: strong delocalization and partly filled energy bands. Such electrons can therefore easily change from one energy state into a slightly different one. Thus, not only do they become delocalized, forming a sea of electrons permeating the lattice, but they are also able to migrate through the lattice when an external electrical field is imposed, leading to electrical conductivity. Without the field there are electrons moving equally in all directions. Under the field some will adjust their state slightly, adopting a different wave vector. Consequently, there will be more moving one way than the other and a net current will result.
The freedom of conduction electrons to migrate also gives metal atoms, or layers of them, the capacity to slide past each other. Locally bonds can easily be broken and replaced by new ones after the deformation. This process does not affect the communal metallic bonding very much. This gives rise to metals' typical characteristic phenomena of malleability and ductility. This is particularly true for pure elements. In the presence of dissolved impurities the defects in the lattice that function as cleavage points may get blocked and the material becomes harder. Gold for example is very soft in pure form (24 kt), which is why for jewelry alloys of 18 kt or lower are preferred.
Metals are typically also good conductors of heat, but the conduction electrons only contribute partly to this phenomenon. Collective (i.e. delocalized) vibrations of the atoms known as phonons that travel through the solid as a wave, contribute strongly.
However, the latter also holds for a substance like diamond. It conducts heat quite well but not electricity. The latter is not a consequence of the fact that delocalization is absent in diamond, but simply that carbon is not electron deficient . The position of carbon in the middle of its period in the Periodic Table makes that there are precisely enough electrons to fill the energy states. Under a field electrons are not able to adopt a different wave vector because there are no empty states to move into. This makes a current impossible in this wide band gap semiconductor. However, as soon as charge carriers are introduced by doping the crystal with a suitable impurity the resulting charge carriers are as mobile as in a metal, though far fewer in number. Even without doping the vibrational motions (the phonons) are delocalized over the crystal explaining the heat conduction. Still the bonding in diamond is better described as covalent than as metallic if only because there is a very strong directional preference for tetrahedral stacking, producing a structure that is extremely hard to deform and by no means close packed.
Clearly, the electron deficiency is an important point in distinguishing metallic from more conventional covalent bonding. Thus, we should amend the expression given above into:
Metallic bonding is an extremely delocalized communal form of electron deficient[7] covalent bonding.
1)-The electromagnetic interaction between delocalized electrons, called conduction electrons, and the metallic nuclei within metals.
-Electropositive elements.
Understood as the sharing of "free" electrons among a lattice of positively-charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals.
2)In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities.
3)Metallic bonding accounts for many physical properties of metals,such as:
-High Melting and Boiling point,(Very strong electrostatic forces of attraction between sea of electrons and the positive metal ions.)
-Strength,
-Malleability and Ductility,(Due to orderly packing of metal atoms.The layers of atom can slide over each other easily whenever a force is applied.)
-Thermal and Electrical conductivity,(They have mobile electrons to carry electricity and heat energy.)
-Opacity, and
-Luster.
4)Although the term metallic bond is often used in contrast to the term covalent bond, it is more preferable to use the term metallic bonding, because this type of bonding is collective in nature and a single "metallic bond" does not exist.
5)The atoms of a metal contribute their valence electrons to form a 'sea' of electrons.
6)The sea of electrons helps to cement the positive ions together to give metals a giant metallic structure.
Physical Properties
1)The physical properties of a substance are determined by its structure and bonding.
STRUCTURE=>PARTICLES=>BONDING BETWEEN PARTICLES=>M.P
Simple Covalent->Simple molecules->Weak Van der waals F.O.A->Low
Macromolecules->Atoms->Strong covalent bonds->High
Metallic->+Ions in "sea" of e- =>Strong metallic Bonds->High
Ionic->Ions->Strong Ionic Bonds->High
The nature of metallic bonding
-The combination of two phenomena gives rise to metallic bonding: delocalization of electrons and the availability of a far larger number of delocalized energy states than of delocalized electrons. The latter could be called electron deficiency.
Delocalization
In 2D
Delocalization -bonding that involves more than one pair of atoms held together by one pair of electrons- is most familiar from the example of benzene C6H6, where six electrons from six carbon atoms are engaged in joint aromatic bonding. However, there are other examples like the three-center two-electron bond, prevalent in boron chemistry. The principle can easily be extended over larger aromatic molecules like naphthalene, anthracene and if the process is taken to its extreme: graphite. The latter is an example of a system delocalized in two dimensions. Interestingly, there is an isoelectronic analog of benzene, B3N3H6 (borazine) for which the same argument holds. It has very similar properties to benzene[5] When extended infinitely hexagonal boron nitride BN is obtained with a structure identical to that of graphite, apart from the alternation between boron and nitrogen in each ring. This material is a semiconductor, exemplifying that delocalization is a necessary but not sufficient requirement for conductivity. Electrical conductivity does occur in graphite, because the π and π*-like bands overlap, making it a semimetal, with partly filled bands, fulfilling the other requirement for conductivity.
In 3D
Metal aromaticity in metal clusters is another example of delocalization, this time often in three-dimensional entities. Metals take the delocalization principle to its extreme and one could say that a crystal of a metal represents a single molecule over which all conduction electrons are delocalized in all three dimensions. This means that inside the metal one can generally not distinguish molecules so that the metallic bonding is neither intra- nor intermolecular. 'Nonmolecular' would perhaps be a better term. Metallic bonding is mostly non-polar, because even in alloys there is little difference among the electronegativities of the atoms participating in the bonding interaction (and in pure elemental metals, none at all). Thus metallic bonding is an extremely delocalized communal form of covalent bonding. In a sense metallic bonding is not a 'new' type of bonding at all therefore and it only describes the bonding as present in a chunk of condensed matter, be it crystalline solid, liquid or even glass. Metallic vapors by contrast are often atomic (Hg) or at times contain molecules like Na2 held together by a more conventional covalent bond. This is why it is not correct to speak of a single 'metallic bond'.
The delocalization is most pronounced for s- and p-electrons. For cesium it is so strong that the electrons are virtually free from the cesium atoms to form a gas only constrained by the surface of the metal. For cesium therefore the picture of Cs+-ions held together by a negatively charged electron gas is not too inaccurate [6]. For other elements the electrons are less free, in that they still experience the potential of the metal atoms, sometimes quite strongly. They require a more intricate quantum mechanical treatment (e.g. tight binding) in which the atoms are viewed as neutral much like the carbon atoms in benzene. For d- and especially f-electrons the delocalization is not strong at all and this explains why these electrons are able to continue behaving as unpaired electrons that retain their spin, adding interesting magnetic properties to these metals.
Electron deficiency and mobility
Metal atoms contain few electrons in their valence shells relative to their periods or energy levels. They are electron deficient elements and the communal sharing does not change that. There remain far more available energy states than there are shared electrons. Both requirements for conductivity are therefore fulfilled: strong delocalization and partly filled energy bands. Such electrons can therefore easily change from one energy state into a slightly different one. Thus, not only do they become delocalized, forming a sea of electrons permeating the lattice, but they are also able to migrate through the lattice when an external electrical field is imposed, leading to electrical conductivity. Without the field there are electrons moving equally in all directions. Under the field some will adjust their state slightly, adopting a different wave vector. Consequently, there will be more moving one way than the other and a net current will result.
The freedom of conduction electrons to migrate also gives metal atoms, or layers of them, the capacity to slide past each other. Locally bonds can easily be broken and replaced by new ones after the deformation. This process does not affect the communal metallic bonding very much. This gives rise to metals' typical characteristic phenomena of malleability and ductility. This is particularly true for pure elements. In the presence of dissolved impurities the defects in the lattice that function as cleavage points may get blocked and the material becomes harder. Gold for example is very soft in pure form (24 kt), which is why for jewelry alloys of 18 kt or lower are preferred.
Metals are typically also good conductors of heat, but the conduction electrons only contribute partly to this phenomenon. Collective (i.e. delocalized) vibrations of the atoms known as phonons that travel through the solid as a wave, contribute strongly.
However, the latter also holds for a substance like diamond. It conducts heat quite well but not electricity. The latter is not a consequence of the fact that delocalization is absent in diamond, but simply that carbon is not electron deficient . The position of carbon in the middle of its period in the Periodic Table makes that there are precisely enough electrons to fill the energy states. Under a field electrons are not able to adopt a different wave vector because there are no empty states to move into. This makes a current impossible in this wide band gap semiconductor. However, as soon as charge carriers are introduced by doping the crystal with a suitable impurity the resulting charge carriers are as mobile as in a metal, though far fewer in number. Even without doping the vibrational motions (the phonons) are delocalized over the crystal explaining the heat conduction. Still the bonding in diamond is better described as covalent than as metallic if only because there is a very strong directional preference for tetrahedral stacking, producing a structure that is extremely hard to deform and by no means close packed.
Clearly, the electron deficiency is an important point in distinguishing metallic from more conventional covalent bonding. Thus, we should amend the expression given above into:
Metallic bonding is an extremely delocalized communal form of electron deficient[7] covalent bonding.
Take a Break.Have a Kit-Kat.
.. I'M GETTING SO FAT
I CAN HARDLY SCRATCH MY OWN BUTT
THESE MORNING WALKS ARE KILLING ME
OK, NOW DON'T MOVE FOR ABOUT A WEEK
SORRY MOM, I'M NEW AT THIS
I hate it when this happens
I REALLY NEED TO GET GOING,
BUT JUSTCAN'T SEEM TO GET MOTIVATED
Hey..Give Me Back my Ball..
I said, Go to sleep
Nice doggie...GOOD boy
A little power nap...
WHAT PETS DO WHEN WE'RE AT WORK
I CAN HARDLY SCRATCH MY OWN BUTT
THESE MORNING WALKS ARE KILLING ME
OK, NOW DON'T MOVE FOR ABOUT A WEEK
SORRY MOM, I'M NEW AT THIS
I hate it when this happens
I REALLY NEED TO GET GOING,
BUT JUSTCAN'T SEEM TO GET MOTIVATED
Hey..Give Me Back my Ball..
I said, Go to sleep
Nice doggie...GOOD boy
A little power nap...
WHAT PETS DO WHEN WE'RE AT WORK
15.4.09
Recaps.
Molecules and Covalent Bonds
-Molecules and compounds
1)When 2 or more atoms are chemically bonded together;a molecule is formed.
2)A molecule is a group of atoms chemically bonded together.
3)Most non-metallic elements are found in the molecular form.
4)If a molecule is made up of atoms of 2 or more different elements,a compound is formed.
5)A compound is a substance made when 2 or more elements are joined together by a chemical reaction.
COVALENT BONDING
1)Covalent bonds are usually formed between atoms of non-metals.
2)A sharing of electrons allow both to achieve a stable noble gas configuration.
3)If 2 atoms share
a)one pair of electrons=>a single bond is formed.
b)two pairs of electrons=>a double bond is formed.
c)three pairs of electrons=>a triple bond is formed.
4)Each atom contributes an equal number of electrons for sharing.
PROPERTIES OF COVALENT COMPOUNDS
1)Covalent compounds have low melting and boiling points
The molecules are held together by very weak van der waals force of attraction.Vey little energy is required to overcome these weak intermolecular forces.
2)Covalent compounds do not conduct electricity.
Molecules in covalent compounds do not carry charges.
3)Covalent compounds can dissolve in organic solvents but not in water.
Covalent molecules are not readily hydrated by water molecules.
-Molecules and compounds
1)When 2 or more atoms are chemically bonded together;a molecule is formed.
2)A molecule is a group of atoms chemically bonded together.
3)Most non-metallic elements are found in the molecular form.
4)If a molecule is made up of atoms of 2 or more different elements,a compound is formed.
5)A compound is a substance made when 2 or more elements are joined together by a chemical reaction.
COVALENT BONDING
1)Covalent bonds are usually formed between atoms of non-metals.
2)A sharing of electrons allow both to achieve a stable noble gas configuration.
3)If 2 atoms share
a)one pair of electrons=>a single bond is formed.
b)two pairs of electrons=>a double bond is formed.
c)three pairs of electrons=>a triple bond is formed.
4)Each atom contributes an equal number of electrons for sharing.
PROPERTIES OF COVALENT COMPOUNDS
1)Covalent compounds have low melting and boiling points
The molecules are held together by very weak van der waals force of attraction.Vey little energy is required to overcome these weak intermolecular forces.
2)Covalent compounds do not conduct electricity.
Molecules in covalent compounds do not carry charges.
3)Covalent compounds can dissolve in organic solvents but not in water.
Covalent molecules are not readily hydrated by water molecules.
Recaps.
Formation Of Ions
1)Ions are charged particles.
2)Ions are formed when atoms lose or gain electrons.
3)Atoms lose or gain electrons in order to achieve a noble gas configuration.
4)Metal atoms lose their valence electrons to achieve a noble gas configuration.A positive ion(Cation) is formed.
For Example:
Group I metals lose one valence electron to form ions with charge+1.
Na(2.8.1)---> Na+(2.8)
K (2.8.8.1)--->K+(2.8.8)
6)Non-metal atom gain electrons to achieve a noble gas configuration .A negative ion(Anion)is formed.
For Example:
Group VI non-metals gain two electrons to form ions with charge-2.
O(2.6)--->O^2-(2.8)
S(2.8.6)-->S^2-(2.8.8)
IONIC BONDING
1)Ionic bonds are formed between metal atoms and non-metal atoms.
2)A transfer of electrons from a metal atom to a non-metal atom enables both to achieve a stable noble gas configuration.
3)The resulting oppositely-charged ions then attract each other via Strong Electrostatic Forces Of Attraction.
Sodium and chlorine are the best examples to illustrate how elements form compounds.
Notice how sodium is just one electron more from being stable? and that chlorine is just one electron less from being stable?
Look below and see that when sodium and chlorine atoms bump together, their outer orbits' electron will react with each other.
The positive chlorine nucleus is stronger, so, it is able to pull the loosely held sodium electron into it's orbits.
The transfer of an electron from the metallic sodium to a non-metallic chlorine rewulted in a new structure for both atoms. Sodium has given up one electron, and chlorine accepted it. Now, the arrangement of the atoms are stable, but it no longer has a structure of a neutral atom.
Both the Na and Cl particles are ions, because they each have an electrical charge. They are attracted to each other because of their opposite charges. This force of attraction between + and - ions is called an ionic bond. If an atoms gains electrons, it has a negative charge and tends to be non-metal. However, if an atom loses electrons, it has a positive charge and are usually metallic. It is the attraction between positive sodium ions (Na+1) and negative chloride ions (Cl-1) that holds the new compound together! Therefore, Sodium chloride is called an ionic compound, because it is made up of ions and held together by ionic bonds.
Properties Of Ionic Compounds
1)Ionic Compounds have high melting and boiling points.
-The oppositely-charged ions are held together by very strong electrostatic forces of attraction which requires a lot of energy to overcome.
-Therefore,ionic compounds are difficult to melt.The ability to withstand high temperatures makes ionic compounds(e.g. Magnesium Oxide) suitable for use as lining in furnaces.
2)Ionic Compounds do not conduct electricity in the solid state;It only conduct electricity when molten or dissolved in water.
-In the solid state,the ions are bonded together by strong electrostatic forces of attraction and can only vibrate about a fixed position.The ions cannot move freely to conduct electricity.
-In the molten or aqueous state,the ions are free to move and carry an electric current.
3)Ionic Compounds readily dissolve in water but do not dissolve in organic solvents.
-Ions are readily hydrated by water but not by organic solvent
`
1)Ions are charged particles.
2)Ions are formed when atoms lose or gain electrons.
3)Atoms lose or gain electrons in order to achieve a noble gas configuration.
4)Metal atoms lose their valence electrons to achieve a noble gas configuration.A positive ion(Cation) is formed.
For Example:
Group I metals lose one valence electron to form ions with charge+1.
Na(2.8.1)---> Na+(2.8)
K (2.8.8.1)--->K+(2.8.8)
6)Non-metal atom gain electrons to achieve a noble gas configuration .A negative ion(Anion)is formed.
For Example:
Group VI non-metals gain two electrons to form ions with charge-2.
O(2.6)--->O^2-(2.8)
S(2.8.6)-->S^2-(2.8.8)
IONIC BONDING
1)Ionic bonds are formed between metal atoms and non-metal atoms.
2)A transfer of electrons from a metal atom to a non-metal atom enables both to achieve a stable noble gas configuration.
3)The resulting oppositely-charged ions then attract each other via Strong Electrostatic Forces Of Attraction.
Sodium and chlorine are the best examples to illustrate how elements form compounds.
Notice how sodium is just one electron more from being stable? and that chlorine is just one electron less from being stable?
Look below and see that when sodium and chlorine atoms bump together, their outer orbits' electron will react with each other.
The positive chlorine nucleus is stronger, so, it is able to pull the loosely held sodium electron into it's orbits.
The transfer of an electron from the metallic sodium to a non-metallic chlorine rewulted in a new structure for both atoms. Sodium has given up one electron, and chlorine accepted it. Now, the arrangement of the atoms are stable, but it no longer has a structure of a neutral atom.
Both the Na and Cl particles are ions, because they each have an electrical charge. They are attracted to each other because of their opposite charges. This force of attraction between + and - ions is called an ionic bond. If an atoms gains electrons, it has a negative charge and tends to be non-metal. However, if an atom loses electrons, it has a positive charge and are usually metallic. It is the attraction between positive sodium ions (Na+1) and negative chloride ions (Cl-1) that holds the new compound together! Therefore, Sodium chloride is called an ionic compound, because it is made up of ions and held together by ionic bonds.
Properties Of Ionic Compounds
1)Ionic Compounds have high melting and boiling points.
-The oppositely-charged ions are held together by very strong electrostatic forces of attraction which requires a lot of energy to overcome.
-Therefore,ionic compounds are difficult to melt.The ability to withstand high temperatures makes ionic compounds(e.g. Magnesium Oxide) suitable for use as lining in furnaces.
2)Ionic Compounds do not conduct electricity in the solid state;It only conduct electricity when molten or dissolved in water.
-In the solid state,the ions are bonded together by strong electrostatic forces of attraction and can only vibrate about a fixed position.The ions cannot move freely to conduct electricity.
-In the molten or aqueous state,the ions are free to move and carry an electric current.
3)Ionic Compounds readily dissolve in water but do not dissolve in organic solvents.
-Ions are readily hydrated by water but not by organic solvent
`
14.4.09
Recaps.
Atomic Structure.
1)Matter is built up from small individual particles called atoms.
2)An atom iis made up 3 basic particles:
-Protons
-Neutrons
-Electrons
3)The properties of these three sub-atomic particles are shown below:
ELEMENTS--->RELATIVE MASS--->CHARGE--->REST MASS--->FOUND IN:
Protons(p)---> 1 ---> +1 ---> 1.67 x 10^-27 kg ---> Nucleus
Neutrons(n)---> 1 ---> 0 ---> 1.67 x 10^-27 kg ---> Nucleus
Electrons(e)--->1/1840---> -1 --->9.11 x 10^-31 kg--> Around Nucleus
4)The protons and neutrons are found in the nuleus of the atoms.These 2 types of particles that make up the nucleus are known as Nucleons.The protons and neutrons make up the dense core of the atom.
5)The electrons move around the nucleus in definite energy levels or electron "shells".
6)Te composition of protons ,neutrons and electrons in an atom can be determined from its proton number(atomic number)and nucleon number(mass number).
7)The proton number is the number of protons in an atom.
8)The nucleon number is the total number of nucleons(p+n)in the nucleus of an atom.
9)Since an atom is electrically neutral,the number of protons is the same as the number of electrons.
Therefore,
Number of neutrons=Nucleon number - Proton number
ELEMENTS
1)An elements is a pure substance that cannot be decomposed into anything simpler by a chemical reaction.
ISOTOPES
1)Isotopes are different atoms of an element with the same proton number but different nucleon numbers.
Isotopes of a particular element have the same:
i)Atomic number(proton number)
ii)Number of electrons in a neutral atom
iii)Electronic Configuration(same number of electrons)
iv)Chemical properties(same number of valence electrons)
Isotopes of a particular element have different:
a)Number of neutrons in an atom
b)Mass numbers
c)Relative isotopic mass
d)Physical properties
ELECTRON CONFIGURATION
1)The way in which the electrons are distributed among the various energy levels or shells.
2)Each energy level can accommodate a definite number of electrons.
3)The energy levels must be filled in order of increasing energy.The first level is filled first before going to the second level and subsequent higher levels.
4)The outermost shell is the highest energy level that is occupied by electrons.
5)Electrons in outermost shell are called valence electron.
6)When the outermost shell is completely filled,the atom would have achieved a stable noble gas configuration.
7)Atoms of other element also try to achieve a noble gas configuration by forming chemical bonds with other atoms.
8)The types of chemical bonds formed between atoms are:
a)Ionic (or electrovalent) bond
b)Simple covalent bond
c)Metallic bond
1)Matter is built up from small individual particles called atoms.
2)An atom iis made up 3 basic particles:
-Protons
-Neutrons
-Electrons
3)The properties of these three sub-atomic particles are shown below:
ELEMENTS--->RELATIVE MASS--->CHARGE--->REST MASS--->FOUND IN:
Protons(p)---> 1 ---> +1 ---> 1.67 x 10^-27 kg ---> Nucleus
Neutrons(n)---> 1 ---> 0 ---> 1.67 x 10^-27 kg ---> Nucleus
Electrons(e)--->1/1840---> -1 --->9.11 x 10^-31 kg--> Around Nucleus
4)The protons and neutrons are found in the nuleus of the atoms.These 2 types of particles that make up the nucleus are known as Nucleons.The protons and neutrons make up the dense core of the atom.
5)The electrons move around the nucleus in definite energy levels or electron "shells".
6)Te composition of protons ,neutrons and electrons in an atom can be determined from its proton number(atomic number)and nucleon number(mass number).
7)The proton number is the number of protons in an atom.
8)The nucleon number is the total number of nucleons(p+n)in the nucleus of an atom.
9)Since an atom is electrically neutral,the number of protons is the same as the number of electrons.
Therefore,
Number of neutrons=Nucleon number - Proton number
ELEMENTS
1)An elements is a pure substance that cannot be decomposed into anything simpler by a chemical reaction.
ISOTOPES
1)Isotopes are different atoms of an element with the same proton number but different nucleon numbers.
Isotopes of a particular element have the same:
i)Atomic number(proton number)
ii)Number of electrons in a neutral atom
iii)Electronic Configuration(same number of electrons)
iv)Chemical properties(same number of valence electrons)
Isotopes of a particular element have different:
a)Number of neutrons in an atom
b)Mass numbers
c)Relative isotopic mass
d)Physical properties
ELECTRON CONFIGURATION
1)The way in which the electrons are distributed among the various energy levels or shells.
2)Each energy level can accommodate a definite number of electrons.
3)The energy levels must be filled in order of increasing energy.The first level is filled first before going to the second level and subsequent higher levels.
4)The outermost shell is the highest energy level that is occupied by electrons.
5)Electrons in outermost shell are called valence electron.
6)When the outermost shell is completely filled,the atom would have achieved a stable noble gas configuration.
7)Atoms of other element also try to achieve a noble gas configuration by forming chemical bonds with other atoms.
8)The types of chemical bonds formed between atoms are:
a)Ionic (or electrovalent) bond
b)Simple covalent bond
c)Metallic bond
9.4.09
.JPG)
What's That?
First Ionisation Energy: The energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.
This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X.
Things to take note about the equation:
The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.
Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).
All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table
The first 20 elements
First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.
These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.
Factors affecting the size of ionisation energy:
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.
The size of that attraction will be governed by:
The charge on the nucleus.
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus.
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
The number of electrons between the outer electrons and the nucleus.
Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!)
If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.
Explaining the pattern in the first few elements
Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).
Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.
Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).
If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3
Talking through the next 17 atoms one at a time would take ages. We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends.
The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.
Explaining the general trend across periods 2 and 3
The general trend is for ionisation energies to increase across a period.
In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons.
The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.
In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.
Why the drop between groups 2 and 3 (Be-B and Mg-Al)?
The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest.
Be 1s22s2 1st I.E. = 900 kJ mol-1
B 1s22s22px1 1st I.E. = 799 kJ mol-1
You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects.
-The increased distance results in a reduced attraction and so a reduced ionisation energy.
-The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.
The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2-level.
Mg 1s22s22p63s2 1st I.E. = 736 kJ mol-1
Al 1s22s22p63s23px1 1st I.E. = 577 kJ mol-1
The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.
Why the drop between groups 5 and 6 (N-O and P-S)?
Once again, you might expect the ionisation energy of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time?
N 1s22s22px12py12pz1 1st I.E. = 1400 kJ mol-1
O 1s22s22px22py12pz1 1st I.E. = 1310 kJ mol-1
The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital.
The difference is that in the oxygen case the electron being removed is one of the 2px2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.
The drop in ionisation energy at sulphur is accounted for in the same way.
Trends in ionisation energy down a group
As you go down a group in the Periodic Table ionisation energies generally fall. You have already seen evidence of this in the fact that the ionisation energies in period 3 are all less than those in period 2.
Taking Group 1 as a typical example:
Why is the sodium value less than that of lithium?
There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.
Li 1s22s1 1st I.E. = 519 kJ mol-1
Na 1s22s22p63s1 1st I.E. = 494 kJ mol-1
Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the centre of 1+.
The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the centre of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy.
Similar explanations hold as you go down the rest of this group - or, indeed, any other group.
Trends in ionisation energy in a transition series
Apart from zinc at the end, the other ionisation energies are all much the same.
All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.
As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.
The rise at zinc is easy to explain.
Cu [Ar]3d104s1 1st I.E. = 745 kJ mol-1
Zn [Ar]3d104s2 1st I.E. = 908 kJ mol-1
In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater. There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this case it obviously isn't enough to outweigh the effect of the extra proton.
Ionisation energies and reactivity
The lower the ionisation energy, the more easily this change happens:
You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form.
The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process.
For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it.
However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are.
The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions.
This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X.
Things to take note about the equation:
The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.
Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).
All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table
The first 20 elements
First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.
These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.
Factors affecting the size of ionisation energy:
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.
The size of that attraction will be governed by:
The charge on the nucleus.
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus.
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
The number of electrons between the outer electrons and the nucleus.
Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!)
If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.
Explaining the pattern in the first few elements
Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).
Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.
Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).
If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3
Talking through the next 17 atoms one at a time would take ages. We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends.
The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.
Explaining the general trend across periods 2 and 3
The general trend is for ionisation energies to increase across a period.
In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons.
The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.
In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.
Why the drop between groups 2 and 3 (Be-B and Mg-Al)?
The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest.
Be 1s22s2 1st I.E. = 900 kJ mol-1
B 1s22s22px1 1st I.E. = 799 kJ mol-1
You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects.
-The increased distance results in a reduced attraction and so a reduced ionisation energy.
-The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.
The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2-level.
Mg 1s22s22p63s2 1st I.E. = 736 kJ mol-1
Al 1s22s22p63s23px1 1st I.E. = 577 kJ mol-1
The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.
Why the drop between groups 5 and 6 (N-O and P-S)?
Once again, you might expect the ionisation energy of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time?
N 1s22s22px12py12pz1 1st I.E. = 1400 kJ mol-1
O 1s22s22px22py12pz1 1st I.E. = 1310 kJ mol-1
The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital.
The difference is that in the oxygen case the electron being removed is one of the 2px2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.
The drop in ionisation energy at sulphur is accounted for in the same way.
Trends in ionisation energy down a group
As you go down a group in the Periodic Table ionisation energies generally fall. You have already seen evidence of this in the fact that the ionisation energies in period 3 are all less than those in period 2.
Taking Group 1 as a typical example:
Why is the sodium value less than that of lithium?
There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.
Li 1s22s1 1st I.E. = 519 kJ mol-1
Na 1s22s22p63s1 1st I.E. = 494 kJ mol-1
Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the centre of 1+.
The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the centre of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy.
Similar explanations hold as you go down the rest of this group - or, indeed, any other group.
Trends in ionisation energy in a transition series
Apart from zinc at the end, the other ionisation energies are all much the same.
All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.
As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.
The rise at zinc is easy to explain.
Cu [Ar]3d104s1 1st I.E. = 745 kJ mol-1
Zn [Ar]3d104s2 1st I.E. = 908 kJ mol-1
In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater. There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this case it obviously isn't enough to outweigh the effect of the extra proton.
Ionisation energies and reactivity
The lower the ionisation energy, the more easily this change happens:
You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form.
The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process.
For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it.
However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are.
The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions.
2.4.09
Ionisation Energy.
Chemistry lessons again,Phew!
Learnt something new:
1)Electronic Configuration(Sub-shell Notation)
2)"Noble Gas" Configuration
For this Configuration thing ,
we have to Remember the 3 RULES of filling up orbitals:
a)Aufbau (building up) principle
-Electrons are placed in the orbital of the lowest energy , then the orbital of the next lowest energy and so on (to give maximum stability to the system).
-The order in which the orbitals are filles:
^
l
l __ __ __ __ __
l 3d
l __
l 4s
l __ __ __
l 3p
l __
l 3s
l __ __ __
l (px)(py)(pz)
l 2p
l __
l 2s
l
l __
1s
The Above Diagrams indicates thats the electron(s) will be filled in the following order:
1s->2s->2p->3s->3p->4s->3d->4p->5s->4d->......
b)Pauli Exclusion Principle
-Each orbital can only hold a maximum of 2 Electrons.
-Both Electrons must be of Opposite spins to each other.
=>WHY? The magnetic attraction which results from their opposite spins can counterbalance the Electrical Repulsion which results from their identical charges.
c)Hund's Rule
-The rules states that in a given energy level,the no. of unpaired electrons should be maximum.
-Pairing of electrons takes place only after all orbitals of the same energy level have been filled with one electron.
=>WHY? By occupying different orbitals, the electrons remain as far away as possible ,thus minimising electron-electron repulsion.
One more IMPORTANT factor to note:
-Electrons occupy the 4s orbital first before filling the 3rd orbitals.
=>WHY?->The 4s is of lower energy than the 3rd orbitals.
Learnt something new:
1)Electronic Configuration(Sub-shell Notation)
2)"Noble Gas" Configuration
For this Configuration thing ,
we have to Remember the 3 RULES of filling up orbitals:
a)Aufbau (building up) principle
-Electrons are placed in the orbital of the lowest energy , then the orbital of the next lowest energy and so on (to give maximum stability to the system).
-The order in which the orbitals are filles:
^
l
l __ __ __ __ __
l 3d
l __
l 4s
l __ __ __
l 3p
l __
l 3s
l __ __ __
l (px)(py)(pz)
l 2p
l __
l 2s
l
l __
1s
The Above Diagrams indicates thats the electron(s) will be filled in the following order:
1s->2s->2p->3s->3p->4s->3d->4p->5s->4d->......
b)Pauli Exclusion Principle
-Each orbital can only hold a maximum of 2 Electrons.
-Both Electrons must be of Opposite spins to each other.
=>WHY? The magnetic attraction which results from their opposite spins can counterbalance the Electrical Repulsion which results from their identical charges.
c)Hund's Rule
-The rules states that in a given energy level,the no. of unpaired electrons should be maximum.
-Pairing of electrons takes place only after all orbitals of the same energy level have been filled with one electron.
=>WHY? By occupying different orbitals, the electrons remain as far away as possible ,thus minimising electron-electron repulsion.
One more IMPORTANT factor to note:
-Electrons occupy the 4s orbital first before filling the 3rd orbitals.
=>WHY?->The 4s is of lower energy than the 3rd orbitals.
1.4.09
April's Fools Day.
Interesting Day.
Friends tried to bluff me!
Friends messaged me!
Friend telling me he likes me!
And well, of course.
I didn't fall for their tricks/pranks.
How smart am I!
Friends tried to bluff me!
Friends messaged me!
Friend telling me he likes me!
And well, of course.
I didn't fall for their tricks/pranks.
How smart am I!
Excel.
I did the online assignment-IT Activity today!
About the Ionisation Energy.
Drew out the tables,
had to refer to the DATA Booklet for informations.
Then based on the data ,plot the Graphs.
3 Graphs in total.
At first i didn't know what to do,then found out that there's video to aid us in doing this activity.
Well,those videos are useful.LOLS
But i don't think i had completed it,because they asked us to compare,highlight and explain which I didn't.
And I've submitted it already.
How Wonderful!:)
About the Ionisation Energy.
Drew out the tables,
had to refer to the DATA Booklet for informations.
Then based on the data ,plot the Graphs.
3 Graphs in total.
At first i didn't know what to do,then found out that there's video to aid us in doing this activity.
Well,those videos are useful.LOLS
But i don't think i had completed it,because they asked us to compare,highlight and explain which I didn't.
And I've submitted it already.
How Wonderful!:)
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