chemical substance
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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
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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.
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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.
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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
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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
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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).
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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.
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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.
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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
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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.
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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.
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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)
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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.