Chemical Substance

1 HBSC2103 V2 CHEMISTRY 1

FACULTY OF EDUCATION AND LANGUAGES

SEMESTER MEI 2013

HBSC 2103 V2

CHEMISTRY 1

NAME : AL YASSER BIN MD ISA

MATRICULATION NO : 82022802512900

IDENTITY CARD NO. : 820228-02-5129

TELEPHONE NO. : 012-5662130

E-MAIL : [email protected]

LEARNING CENTRE : SUNGAI PETANI


TABLE OF CONTENT

CONTENT PAGE NO.

1.0 DEFINITION OF CHEMICAL SUBSTANCE

2.0 STRUCTURE OF A CHEMICAL SUBSTANCE

3.0 TYPE OF BONDING

4.0 PROPERTIES OF CHEMICAL SUBSTANCE

5.0 MOLECULAR AND GIANT

MOLECULAR COMPOUNDS OF CARBON

6.0 CONCLUSION

7.0 REFERENCES

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CHEMICAL SUBSTANCE

1.0 DEFINITION OF CHEMICAL SUBSTANCE

Substance means a chemical element and its compounds. The term substance includes

both substances obtained by a chemical manufacturing process (for example formaldehyde or

methanol) and substances in their natural state (e.g., pure metals like lead, cadmium, or minerals,

etc). The term substance also includes its additives and impurities where these are part of its

manufacturing process, but excludes any solvent which can be separated without affecting the

stability of the substance or changing its composition (McMurry, 1992).

Chemical substances can therefore be either a pure chemical element or a pure chemical

compound. There are, however, some other definitions e.g. “A form of matter that has both

definite composition and distinct properties. A pure substance cannot be separated into simpler

components without chemical change. Physical changes can alter the state of matter but not the

chemical identity of a pure substance”.

2.0 STRUCTURE OF A CHEMICAL SUBSTANCE

Structure describes the arrangement in which atoms are held together in space. There are

two extreme types molecular and giant.

A molecular structure is composed of molecules. Small groups of atoms strongly held

together by covalent bonding. The bonds between molecules are much weaker (usually over 100

times weaker than a covalent bond) and are called intermolecular forces. This makes the structure

weaker. It tends to fall apart more easily. This gives a substance that is easy to melt or boil like

water. The substance may be gas or liquid at room temperature. If it is solid, it is likely to be soft.

Examples of molecules include Cl2, H20, H2SO4 and NH3 (Anthony, 1990)

In a giant structure, the atoms are in a regularly repeating arrangement usually in three

dimensions, often called lattice. This lattice may be held together by any of the three types of

strong bond (ionic, covalent or metallic) and the structure is difficult to break up. This gives a

strong substance that is difficult to melt or boil. The substance will be solid, and probably hard

like SiO2 in sand.


3.0 TYPE OF BONDING

A chemical bond is the physical process responsible for the attractive interactions between

atoms and molecules, and that which confers stability to diatomic and polyatomic chemical

compounds. Chemical bond is the force which bounds two or more ions. Ions are formed due to

the fact that every atom or element needs to get stable by either gaining or losing electron(s). By

losing electron cation is formed and by gaining it anion is formed. And these two are attracted by

each other due to opposite charge and the force acting is chemical bond.

The forces of attraction that hold atoms together are called chemical bonds. The following

is a list of different types of chemical bonds:

Intramolecular Forces (bonds within the molecule)

Intramolecular bonds refers to the forces of attraction that hold atoms together within a molecule.

These types of bonds are considerably stronger than intermolecular bonds. Three types of

intramolecular bonds - covalent, ionic and metallic bonds.

a) Covalent Bonds

Covalent bonding is a form of chemical bonding between two non metallic atoms which is

characterized by the sharing of pairs of electrons between atoms and other covalent bonds. A

covalent bond is formed between two non-metals that have similar electronegativities. Neither

atom is "strong" enough to attract electrons from the other. For stabilization, they share their

electrons from outer molecular orbit with others.

Covalent bonds have a definite and predictable shape and have low melting and boiling

points. They can be easily broken into its primary structure as the atoms are close by to share the

electrons. These are mostly gaseous and even a slight negative or positive charge at opposite ends

of a covalent bond gives them molecular polarity.

b) Ionic Bonds

Ionic bond, also known as electrovalent bond is a type of bond formed from the

electrostatic attraction between oppositely charged ions in a chemical compound. These kinds of

bonds occur mainly between a metallic and a non metallic atom. An ionic bond is formed between

a metal and a non-metal. Non-metals (-ve ion) are "stronger" than the metal (+ve ion) and can get

electrons very easily from the metal. These two opposite ions attract each other and form the ionic

bond.


Ionic bonds normally form crystalline atoms and have higher melting points and boiling

points compared to covalent compounds. These conduct electricity in molten or solution state and

they are extremely polar bonds. Most of them are soluble in water but insoluble in non-polar

solvents. They require much more energy than covalent bond to break the bond between them.

c) Metallic Bonds

Force of attraction operating in a metal that holds the atoms together in a metallic

structure. In metallic bonding, metal atoms form a close-packed, regular arrangement. The atoms

lose their outer-shell electrons to become positive ions. The outer electrons become a ‘sea’ of

mobile electrons surrounding a lattice of positive ions. The lattice is held together by the strong

attractive forces between the mobile electrons and the positive ions.

The properties of metals can be explained in terms of metallic bonding. Metals conduct

electricity as the electrons are free to move. Conduction of heat occurs by vibration of the positive

ions as well as via the mobile electrons. Metals are both ductile and malleable because the

bonding is not broken when metals are deformed; instead, the metal ions slide over each other to

new lattice positions.

Intermolecular Forces (bonds between molecules)

Intermolecular bonds refers to the forces of attraction that hold molecules together. These

bonds are considerably weaker than intramolecular bonds. Compounds with stronger

intermolecular forces will be harder to separate therefore they will have larger melting points and

boiling points. This means compounds with weak intermolecular forces are more likely to be

found in gaseous form, while compounds with stronger molecular forces are more likely to be

found as liquids or solids. Three types of intermolecular bonds - hydrogen bonds, dipole-dipole

attractions and van der Waals’ forces.

a) Hydrogen Bonds

A type of dipole-dipole bond that occurs when hydrogen is bonded to a highly

electronegative atom (O, F, N). A hydrogen bond is the attractive force between the hydrogen

attached to an electronegative atom of one molecule and an electronegative atom of a different

molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial

negative charge. The hydrogen then has the partial positive charge. This type of bonds is the

strongest intermolecular force than either van der Waals’ forces or dipole-dipole interactions since

the hydrogen nucleus is extremely small and positively charged and fluorine, oxygen and nitrogen


being very electronegative so that the electron on the hydrogen atom is strongly attracted to the

fluorine, oxygen or nitrogen atom, leaving a highly localised positive charge on the hydrogen

atom and highly negative localised charge on the fluorine, oxygen or nitrogen atom. This means

the electrostatic attraction between these molecules will be greater than for the polar molecules

that do not have hydrogen covalently bonded to either fluorine, oxygen or nitrogen.

b) Dipole-dipole attractions

Occurs when polar molecules align so that their oppositely charged ends are attract each

other. Dipole-dipole forces are attractive forces between the positive end of one polar molecule

and the negative end of another polar molecule. This bond is stronger than van der Waals’ forces

but they are much weaker than ionic or covalent bonds and have a significant effect only when the

molecules involved are close together (touching or almost touching).

c) Van Der Waals’ Forces

Occurs because of momentary dipole (uneven distributions of electrons in a molecule

causing a partially negative end and a partially positive end). These are the weakest

intermolecular forces. These forces arise between molecules due to attraction of electrons to

nuclei of other molecules or due to momentary dipoles when molecules approach each other.

Dispersion forces increase with increase in number of electrons in a molecule. For example

dispersion forces in SiF4 are higher than CH4 due to the fact that SiF4 has more electrons in it

compared to CH4. Dispersion forces are observed only in non polar molecules such as BF3, CH4,

CO2, C2H4 etc.

4.0 PROPERTIES OF CHEMICAL SUBSTANCE

All substances have properties that we can use to identify them. For example we can

idenify a person by their face, their voice, height, finger prints, DNA etc. The more of these

properties that we can identify, the better we know the person. In a similar way matter has

properties - and there are many of them. There are two basic types of properties that we can

associate with matter. These properties are called Physical properties and Chemical properties:

Physical properties: Properties that do not change the chemical nature of matter

Chemical properties: Properties that do change tha chemical nature of matter


The physical properties (e.g. boiling point, conductivity, strength) of a substance depend on its

structure and type of bonding present. Bonding determines the type of structure (Chandra, 2005).

5.0 MOLECULAR AND GIANT MOLECULAR COMPOUNDS OF CARBON

Ionic compounds have regular structures (giant ionic lattices) in which there are strong

electrostatic forces in all directions between oppositely charged ions. These compounds have high

melting points and high boiling points. Atoms that share electrons can also form giant structures

or macromolecules. Diamond and graphite (forms of carbon) and silicon dioxide (silica) are

examples of giant covalent structures (lattices) of atoms (with very high melting points).

Covalent compounds - simple molecules

Covalent bonds form between non-metal atoms. Each bond consists of a shared pair

of electrons, and is very strong. Covalently bonded substances fall into two main types:

1. simple molecules

2. giant covalent structures.

Simple molecules

In chemistry, the structure of a non-metallic element or compound, in which groups of

atoms are joined together by covalent bonds to form a molecule. The Internet Encyclopaedia

of Science (2012) said that, the molecules of non-metals are often made up of two identical

atoms covalently bonded together, for example: oxygen (O2), hydrogen (H2), and nitrogen

(N2). There are also molecular compounds made up of different non-metal atoms bonded

together, such as water (H2O), ammonia (NH3), and carbon dioxide (CO2). Covalent

molecules that contain only a few atoms are called simple covalent structures. They are made

up of independent molecular units.

The atoms in molecules are held together by strong covalent bonds. These bonds act

only between the atoms within the molecules and so simple molecules have little attraction

for each other. The attraction force between molecules is the weak van der Waals’ force and

these forces are overcome when a molecular substance melts or boils. These means that

substances made of small molecules (eg H2, Cl2, CH4, CO2) have the weakest intermolecular

forces and are gases at room temperature (Cotton, 1999). Larger molecules have stronger

attractions and so may be liquids at room temperature (eg Br2 and C6H14) or solids with low

melting points (eg I2).


A few substances that contain simple covalent molecules are solid at room

temperature. These are molecular solids. Iodine is a molecular solid at room temperature.

Two iodine atoms form a single covalent bond to become an iodine molecule. The solid is

formed because millions of iodine are held together by weak forces of attraction to create a

3D molecular lattice. A molecular solid have low melting and boiling points due to the weak

forces of attraction between the molecules can be broken by small amount of energy. This

means that the molecular solids are usually soft and brittle where they shatter when hit.

Molecular solids are also unable to conduct electricity because there are no free electrons or

ions to carry the charge.

A molecule of carbon dioxide

These contain only a few atoms held together by strong covalent bonds. An example is

carbon dioxide (CO2), the molecules of which contain one atom of carbon bonded with two

atoms of oxygen.

Properties of simple molecular substances

Low melting and boiling points - Substances made up of simple molecules have low

melting points because there are strong bonds between the atoms in the molecule, but

weak bonds holding the molecules together. Therefore, the intermolecular forces

break fairly easily, due to the fact that they are weak and the covalent bonds making

up the molecule do not break because they are strong. This means that, due to the

weak intermolecular forces breaking down easily, simple molecular substances have

low melting and boiling points.

Non-conductive - Substances with a simple molecular structure do not

conduct electricity. This is because they do not have any free electrons or an overall

electric charge.


Covalent bonding - giant covalent structures

Deprez (1988) concluded that, giant covalent structures contain a lot of non-metal

atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into

giant regular lattices - extremely strong structures because of the many bonds involved. The

graphic shows the molecular structure of diamond and graphite: two allotropes of carbon, and

of silica (silicon dioxide).

From left to right - graphite, diamond, silica

Properties of giant covalent structures

Very high melting points - Substances with giant covalent structures have very high

melting points, because a lot of strong covalent bonds must be broken. Graphite, for

example, has a melting point of more than 3,600ºC.

Variable conductivity - Diamond does not conduct electricity. Graphite contains

free electrons, so it does conduct electricity. Silicon is semi-conductive - that is,

midway between non-conductive and conductive.


Graphite

Carbon atoms in one layer Layers of carbon atoms

Graphite is a form of carbon in which the carbon atoms form layers which calls the

flat layer of carbon atoms. These layers can slide over each other, in each layer, each carbon

atom joined to the three other carbon atoms by strong covalent bonds. Layers of carbon atoms

are held by weak van der Waals’ forces and there are no covalent bonds between the layers.

Therefore graphite is much softer than diamond. It is used in pencils, and as a lubricant. It has

very high melting/boiling point because a lot of energy required to break the strong covalent

bonds between the carbon atoms within each layer. Savvatimskiy, A. (2005) mentioned that

graphite is soft and slippery because layers of carbon atoms can slide over each other to the

weak van der Waals’ forces between the layers. Each carbon atoms contributes one outer

shell electron which is not used to form covalent bonds. These electrons can move along the

layers to conduct electricity.

Carbonatoms

Covalentbonds

Van derWaals’ forces


Diamond

Carbon atoms

Diamond is a form of carbon in which each carbon atom is joined to four other carbon

atoms, forming a giant covalent structure. Each carbon atom joined tetrahedral to four other

carbon atoms by strong covalent bonds. Lior Itzhaki (2005) stressed, a lot of energy is

required to break the strong covalent bonds between the carbon atoms. As a result, diamond

is very hard and has a high melting point. It does not conduct electricity because no mobile

charge carries (ions or electrons) present.


Silica

Silica, which is found in sand, has a similar structure to diamond. It contains silicon

and oxygen atoms, instead of carbon atoms. Each silicon atom is covalently bonded to four

oxygen atoms by strong covalent bonds in a tetrahedral structure and each oxygen atom is

covalently bonded to two silicon atoms. It is also hard and has a high melting point because a

lot of energy is required to break the strong covalent bonds between the carbon atoms. No

mobile charge carriers (ions or electrons) present thus silica does not conduct electricity. The

fact that it is a semi-conductor makes it immensely useful in the electronics industry: most

transistors are made of silica.

Buckminsterfullerene

Structure of a buckminsterfullerene molecule - a large ball of 60 atoms

Buckminsterfullerene is yet another allotrope of carbon. It is actually not a giant

covalent structure, but a giant molecule in which the carbon atoms form pentagons and

hexagons - in a similar way to a leather football. It is used in lubricants.


6.0 CONCLUSION

Because of the nature of ionic and covalent bonds, the materials produced by those

bonds tend to have quite different macroscopic properties. The atoms of covalent materials

are bound tightly to each other in stable molecules, but those molecules are generally not very

strongly attracted to other molecules in the material. On the other hand, the atoms (ions) in

ionic materials show strong attractions to other ions in their vicinity. This generally leads to

low melting points for covalent solids, and high melting points for ionic solids.

(2766 words)


7.0 REFERENCES

Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W. and Nichols, Monte C., ed. (1990). "Graphite" (PDF). Handbook of Mineralogy. I (Elements, Sulfides, Sulfosalts). Chantilly, VA, US: Mineralogical Society of America.

Chandra, Sulekh. (2005). Comprehensive Inorganic Chemistry. New York: New Age Publishers.

Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley-Interscience.

Deprez, N.; McLachlan, D. S. (1988) The analysis of the electrical conductivity of graphite conductivity of graphite powders during compaction. Journal of Physics D: Applied Physics, Volume 21, Issue 1, pp. 101-107 (1988).

Lior Itzhaki. (2005). Harder than Diamond: Determining the Cross-Sectional Area and Young's Modulus of Molecular Rods. Rusia: Angewandte Chemie.

McMurry, John E. (1992). Organic Chemistry (3rd ed.). Belmont: Wadsworth

Savvatimskiy, A. (2005). Measurements of the melting point of graphite and the properties of liquid carbon (a review for 1963–2003). Carbon 43 (6): 1115.

The Internet Encyclopedia of Science (2012). Biological Abundance of Elements. Retrieved from http://incomprehensibleuniverse.tumblr.com/post/4291017617/the-internet-encyclopedia-of-science.

http://incomprehensibleuniverse.tumblr.com/post/4291017617/the-internet-encyclopedia-of-science

http://incomprehensibleuniverse.tumblr.com/post/4291017617/the-internet-encyclopedia-of-science

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