1)The modern Periodic Table is arranged in order of increasing atomic number. As one moves from one element to another on the right, one more proton is found in the nucleus, and one more electron is found in the same electron 'shell' (energy level). For this reason, all the elements in Period 3 have the first electron 'shell' (energy level) filled with 2 electrons and the second electron 'shell' (energy level) filled with 8 electrons (the electronic configuration of Neon).
For example,Sodium begins a new electron 'shell' ( 3rdenergy level) with 1 electron, Magnesium has 2 electrons in the third electron 'shell' (energy level), Aluminium has 3 electrons in the third electron 'shell' (energy level) etc, until finally the third electron 'shell' (energy level) is filled with 8 electrons and the stable electronic configuration of the Noble Gas Argon is reached (2,8,8).
2)ATOMIC RADII:
Atomic radius of the elements decrease across the Period from left to right. As we move from left to right across the period one more proton is added to the nucleus of each successive atom, and one more electron is added to the same electron 'shell' (energy level) of each successive atom. The increased positive charge in the nucleus of each successive atom attracts all the electrons in the atom more strongly, so they are drawn in more closely towards the nucleus.
3)1st Ionization Energy :
1st Ionization Energy (the energy required to remove an electron from the gaseous atom) increases across the Period from left to right. The further away from the positively charged nucleus that a negatively charged electron is, the less strongly the electron is attracted to the nucleus and so the more easily that electron can be removed. So, as the atomic radius decreases from left to right across the Period so the 1st Ionization Energy increases.
4)Electronegativity :
Electronegativity(the relative tendency shown by an atom to attract electrons to itself) increases across the Period from left to right. Typically, metals have low electronegativity, little ability to attract electrons, while non-metals have high electronegativity, greater ability to attract electrons. As we move from left to right across the Period, the elements become less metallic in nature (more non-metallic).
2.8.09
Relative trend in melting point of period 3 elements.Explain.
1)MELTING AND BOLING POINT:
a)Melting
When a substance melts, some of the attractive forces holding the particles together are broken or loosened so that the particles can move freely around each other but are still close together. The stronger these forces are, the more energy is needed to overcome them and the higher the melting temperature.
b)Boiling
When a substance boils, most of the remaining attractive forces are broken so the particles can move freely and far apart. The stronger the attractive forces are, the more energy is needed to overcome them and the higher the boiling temperature.
-Sodium, magnesium and aluminium
Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalised electrons. Going from sodium to aluminium:
The charge on the metal ions increases from +1 to +3 (with magnesium at +2),
the number of delocalised electrons increases, so the strength of the metallic bonding increases and
the melting points and boiling points increase.
i)Silicon
Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals). Silicon has giant covalent bonding. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. This extends in three dimensions to form a giant molecule or macromolecule.
Silicon has a very high melting point and boiling point because all the silicon atoms are held together by strong covalent bonds which need a very large amount of energy to be broken.
ii)Phosphorus, sulphur, chlorine and argon
These are all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules, with strong covalent bonds between their atoms. Argon exists as separate atoms (it is monatomic).
Their melting and boiling points are very low because:
When these four substances melt or boil, it is the van der Waals’ forces between the molecules which are broken ...
which are very weak bonds thus little energy is needed to overcome them.
Sulphur has a higher melting point and boiling point than the other three because:
phosphorus exists as P4 molecules,sulphur exists as S8 molecules, chlorine exists as Cl2 molecules and
argon exists individual Ar atoms.
The strength of the van der Waals’ forces decreases as the size of the molecule decreases
so the melting points and boiling points decrease in the order S8 > P4 > Cl2 > Ar.
a)Melting
When a substance melts, some of the attractive forces holding the particles together are broken or loosened so that the particles can move freely around each other but are still close together. The stronger these forces are, the more energy is needed to overcome them and the higher the melting temperature.
b)Boiling
When a substance boils, most of the remaining attractive forces are broken so the particles can move freely and far apart. The stronger the attractive forces are, the more energy is needed to overcome them and the higher the boiling temperature.
-Sodium, magnesium and aluminium
Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalised electrons. Going from sodium to aluminium:
The charge on the metal ions increases from +1 to +3 (with magnesium at +2),
the number of delocalised electrons increases, so the strength of the metallic bonding increases and
the melting points and boiling points increase.
i)Silicon
Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals). Silicon has giant covalent bonding. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. This extends in three dimensions to form a giant molecule or macromolecule.
Silicon has a very high melting point and boiling point because all the silicon atoms are held together by strong covalent bonds which need a very large amount of energy to be broken.
ii)Phosphorus, sulphur, chlorine and argon
These are all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules, with strong covalent bonds between their atoms. Argon exists as separate atoms (it is monatomic).
Their melting and boiling points are very low because:
When these four substances melt or boil, it is the van der Waals’ forces between the molecules which are broken ...
which are very weak bonds thus little energy is needed to overcome them.
Sulphur has a higher melting point and boiling point than the other three because:
phosphorus exists as P4 molecules,sulphur exists as S8 molecules, chlorine exists as Cl2 molecules and
argon exists individual Ar atoms.
The strength of the van der Waals’ forces decreases as the size of the molecule decreases
so the melting points and boiling points decrease in the order S8 > P4 > Cl2 > Ar.
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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
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