Language of chemistry
Chemists use terms and phrases which are not familiar to everyone. It is important for
chemistry students to get themselves familiar with the language used by chemists.
This will help the student while studying the subject. You will be introduced in this
unit to the language used in chemistry. Let us take a look at some of the common
terms often encountered in chemistry.
Atom – An atom is the basic / smallest unit of matter.
Molecule – A group of atoms that are bonded together and represent a basic unit of an
element or a compound.
Element – It is the simplest substance that cannot be further subdivided into anything
simpler by chemical reactions. An element consists of only one type of atoms.
Examples:
Hydrogen
Oxygen
Carbon
Aluminum
Chlorine
Compound – It is a substance that is formed by the chemical combination of two more
elements in a fixed proportion by mass.
Examples:
Calcium carbonate
Ammonia
Sulfuric acid
Carbon dioxide
Sodium chloride
Mixture – It is a substance that is made up of two or more substances put together in
any proportion without any chemical reaction taking place.
Examples:
Blood
Sea water
Ink
Coke
Symbol – It is an alphabet or two which is used to represent an atom of the element.
Examples:
Hydrogen (H)
Oxygen (O)
Carbon (C)
Aluminum (Al)
Chlorine (Cl)
Formula – It is a group of symbols that represents a molecule of the substance.
Examples:
CaCO3
HCl
H2SO4
NH3
NaCl
Chemical equation – A chemical equation is a symbolic representation of a chemical
reaction using symbols and formulae.
Examples:
Carbon + Oxygen  Carbon dioxide
C + O2  CO2
.
Sulfur + Oxygen  Sulfur dioxide
S + O2  SO2
Calcium carbonate + Hydrochloric acid  Calcium chloride + water +
carbon dioxide
CaCO3 + HCl  CaCl2 + H2O + CO2
Radical – It is a group of atoms that behave as a single unit in a number of
compounds.
Examples:
Carbonate - CO3 2Bicarbonate - HCO3
Nitrate - NO3 Sulfate – SO4 2Sulfite – SO3 2Chloride – Cl Ammonium– NH4+
Salt – It is a compound formed when an acid reacts with a base.
Examples:
NaCl sodium chloride
KCl potassium chloride
NH4Cl ammonium chloride
MgCO3 magnesium carbonate
NH4NO3 ammonium nitrate
Writing formulae – Writing a formula for a compound means we should know what
elements make up the compound. For example water is made up of hydrogen and
oxygen in the ratio of 1:2. Therefore the formula of water is
H2O.
That means there are 2 hydrogen atoms combining with 1 oxygen atom. Why is this?
It is because oxygen has a valency of 2 while hydrogen has a valency of 1.
How do we remember this? For now it will help to remember a few valencies.
Monovalent
Divalent
Trivalent
Tetra valent
H - hydrogen
Ca - calcium
B - boron
C - carbon
Na - sodium
Mg - magnesium
Al - aluminium
Si - silicon
K - potassium
Ba - barium
N - nitrogen
Ge - germanium
Cl - chlorine
O - oxygen
P - phosphorus
Sn - tin
Br - bromine
S - sulfur
Pb - lead
Mono - valent elements all have + or - 1 charge in their ionic state. [An ion is a
charged atom.] Similarly a divalent element has + or – 2 charge in their ionic state.
A trivalent element has + or – 3 charge in their ionic state, and a tetravalent element
will have + or – 4 in their ionic state.
A compound like sodium chloride has the formula NaCl, because it is made up of two
ions Na+ and Cl – which combine in the ratio of 1:1.
But Calcium chloride will have the formula CaCl2. This is because calcium ion is Ca2+
and chloride ion is Cl -. To satisfy the 2+ charge on the calcium ion two chloride ions
need to combine with one calcium ion. This is why the formula becomes CaCl2. Let
us look at aluminium oxide. It is made of Al3+ and O2-. The formula is Al2O3 because
2 Al3+ and 3O2- satisfy the need to equal the charges.
Worksheet 1 – Formulae
Write the answers for these questions in your note book.
1. Write the formula for each of the following compounds:
i) potassium nitrate
ii) boron trichloride
iii) ammonium chloride
iv) calcium chloride
v) barium nitrate
vi) magnesium carbonate
v) aluminium chloride
2. Write the names of the following substances:
i) NH4NO3
ii) P2O5
iii) NaOH
iv) HNO3
v) KMnO4
vi) H3PO4
v) K2Cr2O7
3. Use the terms one from each group to make as many names of chemical
compounds as possible:
Group 1
Group 2
Compound
Barium
sulfate
-----------------------Potassium
carbonate
-----------------------Calcium
nitrate
-----------------------Ammonium
phosphate
-----------------------Sodium
hydroxide
-----------------------Aluminium
chloride
-----------------------Lithium
sulfite
-----------------------Sulfur
oxide
-----------------------Carbon
nitrite
-----------------------4. Write the formula for each of the chemical compounds that you worked out in
previous question.
Worksheet 2 – Balancing chemical equations
Balancing chemical equations – Chemical equations must be balanced as the atoms in
substances are not created new or destroyed. They are only re-arranged in different
ways. Therefore the total number of atoms of the different elements on the reactant
and product sides must be equal.
Examples:
H2 + O2  H2 O
This equation is not balanced because there are 2 oxygen atoms on the
reactant side and only 1 on the product side. To balance the number of atoms on both
sides of the chemical equation we must add 2 before H2O to get:
H2 + O2  2H2O
Similarly balance the following equations:
Mg + O2  MgO
MgCO3 + HCl  MgCl2 + H2O + CO2
Cu + AgNO3  Cu (NO3) 2 + Ag
Na2O + H2O
 NaOH
PCl3 + H2O
 H3PO3 +
+
O2

P2 O5
Si +
Cl2

Si Cl4
P
MgO + HCl 
MgCl2 + H2O
P4O10 + H2O

H3PO4
Cl2O +

HClO
H2 O
Al2O3 + H2SO4 
Al2(SO4) 3
Ca(OH)2 + NH4Cl 
CO2
+
H2 O
HCl

CaCl2 +
C6H12O6 +
+
H2 O
H2O + NH3
O2
States of matter All matter in the universe exists in any of four different states. They are: 1. Solid state 2. Liquid state 3. Gaseous state 4. Plasma state Among these the first three states are found on earth naturally. The plasma state is present in the stars. Plasma state is similar to gaseous state but in which some of the particles are in an ionic state – (positive ions and electrons). We will consider only the first three states of matter. Particles in solids, liquids and gases Particles in a solid – They are packed very closely together. There are strong forces holding the particles together. They are unable to move from one place to another. They can only vibrate. Increasing the temperature increases their kinetic energy (K.E.). As a result, they are able to vibrate faster. Particles in a liquid – Liquid particles are loosely packed and are able to move by sliding over each other. They are close together and their intermolecular forces are weak compared to solids. Particles in a gas – The particles in a gas are far apart and have negligible intermolecular forces. They have high kinetic energy and are free to move randomly. Therefore, they can occupy any volume. They usually occupy the volume of the container. The three states of matter are interconvertible. In terms of kinetic theory of matter, solids have least K.E while the gases have maximum K.E. Molecules have very high Molecules have weak Intermolecular force is intermolecular force. intermolecular force. negligible. Kinetic energy of the K.E of the molecules is more K.E of the molecules is molecules is the least. than in solid state. the greatest. When a substance in its gaseous state is cooled, the molecules lose their K.E and slow down. As their K.E keeps on decreasing, the molecules come close to each other and then their intermolecular force increases. At this point, the gas changes into liquid state. Similarly, when a liquid is cooled the molecules lose K.E further and change into solid state. Heating a substance causes the same changes in reverse. Interconversion of states of matter Diffusion It is the movement of molecules from a region of higher concentration to a region of lower concentration until their concentrations become equal. Many natural processes involve diffusion, for example: -­‐ Spreading of perfume molecules in air Absorption of water by root hairs in plants Respiration in animals Absorption of food from the intestines into the blood Rate of Diffusion The rate of diffusion of a substance depends upon its mass. Heavier substances diffuse slowly while lighter substances diffuse fast. Hence we can state that diffusion of a substance is inversely proportional to its mass. Osmosis Osmosis is the movement of water molecules from a weaker solution to a stronger solution across a selectively permeable membrane. Osmosis is nothing but diffusion. It is the movement of water molecules from a region of higher concentration (weak solution) to a region of lower concentration (stronger solution). Purity of substances The purity of a substance can be determined by measuring the melting, freeing or boiling points. It can also be found by measuring the density of the substance. The presence of the impurity elevates the boiling point and lowers the freezing point. The presence of an impurity also affects the density of the substance. Separation Techniques The differences in the physical and chemical properties of substances are taken advantage of to separate the substances from a mixture. The different techniques used are -­‐ • Filtration • Evaporation • Distillation • Fractional Distillation • Sublimation • Crystallization • Magnetic Separation • Paper Chromatography Chromatography Chromatography is a technique used to separate the components of a mixture. Paper Chromatography Among the various types of chromatography, paper chromatography is the simplest. Special type of paper is used to separate the components in a mixture. These components of the mixture are adsorbed on the surface of the paper and are carried along by the solvent molecules. The lighter components are carried faster than the heavier ones. In this way, in a given length of time, the components travel different distances on the strips of chromatography paper. The ratio of the distance travelled by the substance to the distance travelled by the solvent is known as value. The value of different components is different. The value of a compound is always the same for a given solvent, under the same conditions. Many components in a mixture are colourless and are difficult to identify on the paper. Therefore, specific chemicals called locating agents are sprayed onto the chromatograph paper. The colourless components appear coloured and are then identified. Distillation Distillation: -­‐ distillation is evaporation followed by condensation. This process is used to get pure liquids from a mixture. Eg. We can get pure water from seawater by distillation. Fractional Distillation Fractional Distillation: -­‐ It is a process of separating different liquids from a mixture of liquids by boiling and condensing the components at their respective boiling points. Eg. Ethanol from a mixture of ethanol and water. Petrol, Diesel, and other substances from petroleum are obtained in industry by a modified fractionating column. The petroleum or crude oil as is commonly known is vaporised and fed into the tall fractionating columns. The vapours cool as they rise upward and condense at different levels. These condensed liquids called fractions are collected through separate pipes as shown in the diagram. Worksheet 3 – Separation of mixtures 1. Which stages occur in distillation?
A condensation then evaporation
B condensation then filtration
C evaporation then condensation
D filtration then evaporation
2. Some paraffin is contaminated with soot (carbon). The soot is removed as shown.
Which method is used to remove the soot?
A cracking
B crystallization
C diffusion
D filtration
3. A shirt is stained with red ink from a pen. The shirt is left to soak in a bowl of water.
Which process causes the red colour to spread?
A diffusion
B evaporation
C melting
D neutralisation
4. A solution of copper sulphate was made by reacting excess copper oxide with dilute sulphuric acid. The unreacted copper oxide was removed by filtration. Draw a labelled diagram to show how the mixture can be filtered. 5. The apparatus below was used to separate ethanol from water. i) Complete the empty boxes to name the pieces of apparatus.
ii) Indicate by an arrow where heat is applied.
iii) Name this separation process. Atomic Structure Atoms are mainly composed of three types of particles – electrons, protons and neutrons. Electrons are negatively charged particles while protons are positively charged particles. Neutrons are neutral particles with no charge. Protons and neutrons are present in the nucleus of the atom. The electrons revolve around the nucleus at fixed distances called orbits or shells. The number of protons in the nucleus of an atom is its atomic number. The number of protons and the number of neutrons (number of nucleons) is the mass number of the element. Structure of atoms Hydrogen atom Helium atom Lithium atom Beryllium atom Boron atom Carbon atom Nitrogen atom Oxygen atom Electron Configuration It is the arrangement of electrons in the different shells of an atom. The arrangement of electrons in each shell of the atom can be calculated by using the formula 2n2 where n represents the number of the shell 1, 2, 3…etc. The number of electrons that fill each of the shells can be calculated as shown below 1st shell, n = 1 therefore 2 x 12 = 2 2nd shell, n = 2 therefore 2 x 22 = 8 3rd shell, n = 3 therefore 2 x 32 = 18 4th shell, n = 4 therefore 2 x 42 = 32 And so on. However, the last shell cannot hold more than eight electrons and the last but one (penultimate) shell cannot hold more than 18 electrons. So as these rules are followed we notice that the electron arrangement for the elements works out quite simply as given below. 11Na -­‐ 2, 8, 1 20Ca – 2, 8, 8, 2 Applying the rule stated above 2 x n2 we must have 10 electrons in the 3rd shell. However, since the last shell cannot hold more than 8 electrons the electron arrangement becomes simply 2, 8, 8, 2 for calcium atom. Write the electron arrangement for the first 20 elements H to Ca. Use the periodic table given to find out the atomic number of the elements. Worksheet 4 – Drawing atomic structures Using the examples shown above, draw the atomic structures of the following; 7N14 8 O16 9F
19
10Ne20 11Na
23
24
12Mg 27
14Si
31
15P 16S
35
40
18Ar 39
19K 20Ca
13Al
17Cl
28
32
40
Relative Atomic Mass (RAM) The relative atomic mass of an element is the number of times one atom of the element is heavier than 1/12th the mass of a carbon – 12 atom. Since it is difficult to isolate and measure the mass of an individual atom chemists measure the mass of the different elements by comparing their mass to that of a standard. The standard chosen is the carbon isotope C-­‐12. Carbon is assigned the mass 12.00 atomic units (a.m.u) and the atomic mass of every other element is compared to it. The instrument used to measure the relative atomic mass of an element is the ‘Mass spectrometer’ The Mass Spectrometer The above diagram gives an idea of how a mass spectrometer works. It is not necessary to know the working of the instrument in detail. At this level it is important to know that the mass of an element in atomic units (a.m.u.) is measured relative to the mass of the carbon atom. Looking at the periodic table we come across atomic masses of elements like chlorine having fractional atomic units. This is because of the presence of isotopes. Isotopes – These are atoms of the same element having the same atomic number (number of protons) but different mass number (number of neutrons) Isotopes are atoms of the same element with different number of neutrons. For example hydrogen has three isotopes 1H1, 1H2, 1H3. All the three isotopes have one proton in the nucleus but have different number of neutrons. Similarly, carbon has three isotopes 6C12, 6C13, and 6C14. Do some research on the internet and find out about the other isotopes. Relative Molecular Mass (RMM) The relative molecular mass of a substance is the number of times one molecule of the substance is heavier than 1/12th the mass of a carbon – 12 atom. Similar to the measurement of the atomic mass, the molecular mass is measured using a mass spectrometer. We can compute the relative mass number of the substance as shown in the example below. Examples: -­‐ RMM of H2 = 1 + 1 = 2 N2 = 14 + 14 = 28 CO2 = 12 + 16 + 16 = 44 SO2 = 32 + 16 + 16 = 64 HCl = 1 + 35.5 = 36.5 H2SO4 = 1 + 1 + 32 + 16 + 16 + 16 + 16 = 98 HNO3 = 1 + 14 + 16 + 16 + 16 = 63 NH4Cl = 14 + 1 + 1 + 1 + 1 + 35.5 = 53.5 CaCO3 = 40 + 12 + 16 + 16 + 16 = 100 Use the periodic table to calculate the RMM of the following substances; CuSO4 K2Cr2O7 KMnO4 H3PO4
FeSO4 C2H5OH Periodic Table The periodic table is an arrangement of the elements in order of their atomic or proton number. Elements with similar chemical properties are placed in the same vertical column. These vertical columns are known as Groups. The periodic table also shows horizontal rows. These horizontal rows are called periods. The physical and chemical properties of the elements show variation down the groups as well as across the periods. • Going down a group, the metallic nature of the elements increase • Going across a period, from left to right, the elements change from metallic to non – metallic character • For metallic elements (Groups 1, 2 and 3), the reactivity increases down the group • For non – metallic elements (Groups 5, 6 and 7), reactivity decreases down the group • For metals, the melting and boiling points decrease down the group. However, for non – metals, the melting and boiling points increase down the group Group I elements (Alkali metals) Li, Na, K, Rb, Cs and Fr All these elements have one valence electron in their last shell. Therefore, they all show valency +1. Physical Properties  They are all very soft metals which can be easily cut  Freshly cut metals have a silvery, shiny surface  They are all good conductors of heat and electricity  They all have very low density Variation of physical properties Lithium Sodium Density Melting point Boiling point Potassium increases decreases decreases Rubidium Caesium Variation of chemical properties Lithium Sodium Reactivity Potassium increases Rubidium Caesium Chemical Properties The Alkali metals are the most reactive. Therefore, they are kept under oil to prevent reaction with oxygen and moisture in the air. They all react violently with cold water. 2Li + 2H2O 2LiOH + H2 2Na + 2H2O 2NaOH + H2 2K + 2H2O 2KOH + H2 Group VII elements (The Halogens) F, Cl, Br, I and At The Halogens are all covalent, diatomic molecules. ie. They have two atoms per molecule. They all have seven valence electrons and therefore show a valency of -­‐1. Physical Properties These elements are all coloured substances. Fluorine is a pale yellow gas Chlorine is a greenish – yellow gas Bromine is a red – brown liquid Iodine is a greyish – black crystalline solid The melting and boiling points of the Halogens increase down the group. Variation of physical properties Fluorine Chlorine Density Melting point Boiling point Bromine increases increases increases Iodine Variation of chemical properties Fluorine Chlorine Reactivity Bromine decreases Iodine Chemical Properties The reactivity of the elements decreases down the group. The Halogens react with metals to form salts. Example: 2Na + Cl2 2NaCl 2K + Br2 2KBr The element higher up in the group can displace the element below it from its salt solution. Example: F2 + 2KCl Cl2 + 2KBr Br2 + 2KI 2KF + Cl2 2KCl + Br2 2KBr + I2 Transition elements These elements are metals which are placed between Group II and Group III in the periodic table. They form three series. The 1st one starts from Scandium to Zinc. Properties  Transition metals form coloured compounds.  They have variable valancies.  They are all metals with very high melting and boiling points.  Transition metals and their compounds are good catalysts. 0 group elements (Noble gases) He, Ne, Ar, Kr, Xe and Rn The zero group elements are known as noble gases or inert gases. They all have eight electrons in the valance shell except Helium which has two electrons. Since their valence shells are full, they are very unreactive. They neither lose, gain or share electrons. Therefore, their valency is zero. Worksheet 5 – Periodic Table
1
Which statement is correct, about the Periodic Table? Tick its box.
A
It shows 8 periods.
B
The elements are arranged in order of their nucleon numbers.
C
The number of electrons increases by 1 from one element to the next, across a period.
D
Reactivity increases as you move down each group in the table, except for Group 0.
2
This shows the main groups in the first four periods of the Periodic Table.
Statements a – n below describe different elements. Your task is to write the letters a – n,
and the symbols for the corresponding elements, in the correct places in the table.
a
a solid in Period 2 which is quite soft, floats on water, and reacts steadily with it
b
the most reactive non-metal element
c
a green gas that forms diatomic molecules
d
a liquid that does not conduct electricity
e
the element that has two forms, graphite and diamond
f
a gas used to provide an inert atmosphere, for example in light bulbs
g
a colourless gas in which many substances burn readily
h
of all the element in these four periods, it reacts the most violently with water
i
the flammable gas produced when metals react with acid
j
an element which, in ribbon form, burns with a white light, forming ions with a charge of 2+
k
an unreactive gas that makes up most of the air around us
l
one compound of this metal is called limestone
m
a noble gas in Period 4
n
an alkali metal in Period 3
3
The groups in the Periodic Table show trends in their properties.
a Write in two properties that show trends, for Groups I and VII:
i Th
i This increases down the group:
i
………………………………..
This increases down the group:
………………………………
ii This increases up the group:
ii This increases up the group:
…………………………………………………
………………………………………
b The next element in Group I is caesium. How will it compare to the elements above it, for
those two properties?
..............................................................................................................................................
c The next element in Group VII is astatine. How will it compare to the elements above it, for
those two properties?
....................................................................................................................................
4
This table shows observations for reactions between halogens and halide ions.
When this …
chlorine (Cl )
2
Is added to a colourless solution containing …
chloride ions (Cl – ) bromide ions (Br – )
iodide ions (I – )
there is no change
the solution turns
the solution turns
orange
red-brown
bromine (Br )
there is no change
there is no change
the solution turns
red-brown
iodine (I )
there is no change
there is no change
there is no change
2
2
a
i
What is responsible for the orange colour?
ii What is responsible for the red-brown colour?
b
Explain how these results show that:
i
chlorine is the most reactive of those three halogens
ii iodine is the least reactive of those three halogens
Chemical Bonding 1.
2.
3.
4.
Ionic Bond Covalent Bond Dative Bond Metallic Bond 1) Ionic Bond: -­‐ It is a bond formed between two atoms by the transfer of electrons. Eg. NaCl -­‐ Na has 1 electron in its valence shell (outermost shell). During its reaction with Chlorine, the Na atom loses the single valence electron to form a Sodium ion (Na+). The Chlorine atom takes this single electron from the Sodium to form a Chloride ion (Cl-­‐). Chlorine takes this electron from Sodium because it has 7 valence electrons and needs one more to become stable. The positive Na+ ion and the negative Cl-­‐ ion attract each other and form an ionic bond. Dot and Cross Diagrams NaCl CaO When a metallic element and a non -­‐ metallic element react to form a compound, the bond formed is ionic. This is because metal atoms have a tendency to lose electrons while non -­‐ metals have a tendency to gain electrons. 2) Covalent Bond: -­‐ It is a bond formed between two atoms by the sharing of a pair of valence electrons. H2 O2 Whenever a non – metal atom combines with another non – metal atom, the bond formed is always covalent. When a metalloid atom combines with a non – metal atom, the bond formed is also covalent. This is because non – metals have a tendency to take electrons and achieve stability. Therefore, they form binds by sharing their valence electrons. Differences between ionic and covalent compounds Sl no. 1. 2. 3. 4. Ionic Ionic compounds have high M.P and B.P They are solids at room temperature (25˚C) They are crystalline in nature Covalent Covalent compounds have low M.P and B.P They are liquids or gases at room temperature Their solids are soft non – crystalline substances They do not conduct electricity They conduct electricity in molten state or in aqueous solution 5. Ionic compounds are soluble in They are insoluble in polar solvents polar solvents like water These are generalized properties and there are always exceptions. Polar Covalent Bonds If there is a covalent bond between and electropositive and an electronegative element, the bond will be polar. This is because the shared pair of electrons is closer to the electronegative atom. Eg. HCl Some elements have a greater tendency to attract electrons than others. Such elements are referred to as highly electronegative elements. In the HCl molecule the Cl atom has a fractional negative ( -­‐) charge. This is because the hydrogen atom has a lower electronegativity while the chlorine atom has a high electronegativity. The shared pair of electrons that form the covalent bond are drawn closer to the chlorine atom giving it a fractional negative charge. The opposite end of the molecule has a fractional positive ( +) charge, making the molecule polar in nature. If there is an electronegative difference between the atoms that bond to each other, the covalent bond becomes polar. Water is another example of polar covalent compound 3) Dative Bond: -­‐ It is also called co-­‐ordinate covalent bond. A dative bond is a bond formed by the sharing of a pair of electrons between two atoms in which one of them only donates a pair of electrons. Example: Formation of ammonia ion (NH4+) NH3 + H+ NH4+ Ammonia Ammonium ion Formation of hydronium ion (H30) H20 + H+ H3O+ Water Hydronium ion 4) Metallic Bond: -­‐ In metals, the atoms are held by strong metallic bonds. The valence electrons are delocalized and therefore it appears that positively charged atoms are scattered in a sea of electrons. The attractive force between the electrical field created by the electrons and the positively charged atoms is the metallic bond. Metallic bonds account for the unique properties of metals: 1. Metals have high melting and boiling points. 2. They are malleable and ductile. 3. They are good conductors of heat and electricity. 4. They have lustre. 5. They have high density. Macromolecular Structure Some substances are formed by multiple bonds of elements. These substances have very large structures unlike other substances that have small simple molecules. Such substances are known as macromolecules. Diamond, graphite, silica and fullerenes are some of the examples. . Diamond Diamond has a large macromolecular structure made up of pure Carbon atoms. The Carbon atoms bond with each other repeatedly to give macromolecular structure like this. Graphite In Graphite, the Carbon atoms show three bonds with one electron free per atom. The Carbon atoms bond with each other forming hexagonal rings. These hexagonal rings are arranged in layers. These layers are held by weak van der Waals’ forces. Comparison of Diamond and Graphite Structures Sl/no. Diamond Graphite 1. Hard, Crystalline solid Soft, Crystalline solid 2. Poor conductor of electricity Good conductor of electricity 3. Shows tetrahedral structure Shows hexagonal rings structure 4. High density Low density 5. Has lustre Dull appearance Though Diamond and Graphite look different, they are the allotropes of Carbon . Silicon Dioxide (SiO2) Silica or SiO2 or Quartz is a crystalline form if Silicon and Oxygen atoms bonded together. Silicon like Carbon shows valency 4 and bonds with four Oxygen atoms forming SiO2. The structure is very similar to a tetrahedron. It is not as hard as Diamond but compared to many substances it is quite hard. By repeated bonding, SiO2 forms a large macromolecular structure. The covalent bond in the macromolecules is quite strong. All macromolecules have high melting and boiling points because during melting a very large number of bonds have to be loosened. Buckminster fullerenes [Bucky balls] C 60 C70 Buckminster fullerenes are spherical structures composed of carbon atoms covalently bonded together. There are other types of fullerenes containing more number of carbon atoms. They are C60, C70, C72 and C84. These fullerenes are present in chimney soot. When carbon is vaporized and condensed they form unique spherical structures fullerenes. Recently another form of these carbon allotropes called ‘graphene’ was discovered. They are thin sheets of Carbon. Graphene structure Worksheet - 6 Bonding
1.
This is about the bonding in molecules of water, methane, and hydrogen
chloride.
a) First, draw hydrogen atoms in the boxes, to complete the structures of the
molecules.
b) Then use • and x to show their bonding. (Use x for an electron from
hydrogen.)
water
2.
methane
hydrogen chloride
This diagram shows the structure
of a common substance.
a) Extend the structure to the right, by adding
four more ions.
b)
i) Name the substance that has this structure.
……………………………………….
ii) Which type of bonding does it have?
...............................................................
iii) Which word describes the structure, giant or molecular?
................................
c) From the structure, it is possible to predict many properties of the substance.
Underline the most likely property for the solid, in each pair below.
i
ii
iii
solubility in water
melting point / ˚C
electrical conductivity
soluble / insoluble
59 / 801
good / poor
d
Complete the diagrams for the ions in the structure, to show their
electron arrangement.
Show the missing electron shells. (The dark circles show the nucleii.)
e) Explain
how
electrons are
transferred,
when the ions in d
are formed
from their atoms.
…………………………………………………………………………………………
…………………………………………………………………………………………
…………………………………………………………………………………………
3.
These diagrams show
part of the structures of
diamond and graphite.
………………….
………………………..
a) Which do these structures represent, elements or compounds?
.........................................
b) Fill in the three missing labels, for the atom and two structures.
c) Describe the differences in the bonding and structure of graphite and
diamond.
Bonding
…………………………………………………………………………………………
…………………………………………………………………………………………
Structure
…………………………………………………………………………………………
…………………………………………………………………………………………
d) i) One of the two substances is very hard, and the other is soft. Explain this
difference.
…………………………………………………………………………………………
…………………………………………………………………………………………
…………………………………………………………………………………………
…………………………………………………………………………………………
ii) Which substance is therefore used in cutting tools, and which is used as a
lubricant?
cutting tools: …………………………. lubricant: ……………………………………
e) One substance is an insulator, and the other is a good conductor of
electricity. Explain this difference.
…………………………………………………………………………………………
……………….…………………………………………………………………………
………………..………………………………………………………………..……… Quantitative Chemistry Mole:-­‐ It is a fixed quantity of substance that contains Avogadro number of particles (6.02 x 1023) Gram atomic/molecular mass of any substance is one mole of that substance. If an element exists in monoatomic state then the atomic mass of the element is equal to one mole. Substance RMM/RAM One mole No. of particles CO2 44 44 g 6.023 x 1023 molecules NH3 17 17 g 6.023 x 1023 molecules N2 28 28 g 6.023 x 1023 molecules Ca 40 40 g 6.023 x 1023 atoms He 4 4 g 6.023 x 1023 atoms Empirical formula-­‐It shows the simplest ratio in which atoms combine. Molar volume-­‐ It is the volume occupied by 1 mole of any gas at S.T.P. (standard temperature and pressure). One mole of any gas occupies 22.4 dm-­‐3 at S.T.P. [Standard temperature = 00C; Standard pressure = 1 atm] According to Boyle’s Law, the volume of a gas is inversely proportional to its pressure at constant temperature. According to Charles’ Law, volume of a gas is directly proportional to its temperature on a Kelvin scale. Combining the two laws, we get Or, ie This constant is the Gas Constant (R). We can rearrange this and write PV = RT for one mole of any gas. Or PV = nRT for ‘n moles’ of the gas. This is the general gas equation. Example-­‐ What is the volume occupied by 0.25 moles of a gas at S.T.P.? 1.0 mol of the gas will occupy 22.4 dm-­‐3 Therefore, 0.25 mol of the gas will occupy 0.25 x 22.4 = 6.0 dm-­‐3. Now solve the following problems-­‐ 1. What is the volume of 16g of oxygen at S.T.P.? 2. How much volume is occupied by 1.7g of NH3 at S.T.P.? 3. If 6.0g of carbon burns completely in air, what volume of CO2 is produced at S.T.P.? Percentage purity – The purity of a substance can be expressed as percentage. Calculating the percentage purity of a substance; Example-­‐ 10g of chalk was analysed and found to contain 9.5g of CaCO3. What is it’s percentage purity? 9.5g x 100 = 95% 10g Now work out the following problems-­‐ 1. Analysis of 80g of Asprin showed that it contained 79.8g of the pure substance. What is the percentage purity? 2. Salt obtained from the sea was found to contain 86% NaCl. How much NaCl can be obtained from 200g of this salt? 3. A 5.0g sample of dry ice was found to contain 4.4g of pure carbon dioxide. What is the percentage purity of the dry ice? Worksheet - 7 Quantitative chemistry
1
Masses from moles, and moles from masses
Fill in the calculation triangle below, for moles of solute, volume, and concentration. Then
complete the table.
Ar or Mr
substance
number
of moles
Cu
2
Mg
0.5
Cl2
mass (g)
35.5
H2
8
P4
2
O3
1.6
H2 O
54
CO2
0.4
NH3
8.5
CaCO3
2.
Ar values: H =
1, C = 12, N =
14, O = 16,
Mg = 24, P =
31, Cl = 35.5,
Ca = 40, Cu =
64)
100
Masses and equations
a ) What mass of iron(III) oxide is needed to produce 100 g of iron, in the blast furnace?
Equation: Fe2O3 (s) + 3CO (g)
2Fe (s) + 3CO2 (g)
(Ar: O = 16, Fe = 56)
b ) 0.05 moles of aluminium is reacted with 26 g of iodine. Which one is the limiting reagent?
Equation: 2Al (s) + 3I2 (s)
2AlI3 (s)
(Ar: Al = 27, I = 127)
26 g of I2 is ……….. moles of I2.
From the equation, this will react with………… moles of Al.
c ) 6.21 g of lead (Pb) are heated in oxygen and give 6.85 g of a lead oxide.
What is the equation for the reaction?
(Ar: O = 16, Pb = 207)
The mass of oxygen that took part in the reaction was ……………. g, which is …………….
moles of O2. The number of moles of Pb in 6.21 g of lead is …………….
So the balanced equation is:
…………………………………………………………………………
3
Calculations involving solutions—The concentration of a solution is the amount of solute,
in grams or moles, that is dissolved in 1 dm3 of solution.
volume
of solution
solute
3
Sodium
chloride
1 dm
hydrochloric
acid
100 cm
Sodium
hydroxide
2 dm
sulfuric acid
concentration
of solution (mol / dm3)
2
3
0.5
3
250 cm
moles of solute in it
1
3
0.5
ammonium
nitrate
copper(II)
sulfate
2
0.3
0.25
0.75
Fill in the calculation triangle below, for moles of solute, volume, and concentration. Then
complete the table.
4
Concentration and equations
a) 25 cm3 of 0.2 mol / dm3 sodium hydroxide (NaOH) neutralises 10 cm3 of dilute sulfuric
acid (H2SO4). What is the concentration of the sulfuric acid?
Equation: 2NaOH (aq) + H2SO4 (aq)
Na2SO4 (aq) + H2O (l)
…………….. moles of NaOH are used, so they neutralise …………….. moles of H2SO4.
So the concentration of the sulfuric acid is …………….. mol / dm3.
b) What mass of magnesium will react with 250 cm3 of 2 mol/dm3 hydrochloric acid?
Equation: Mg (s) + 2HCl (aq)
MgCl2 (aq) + H2 (g)
(Ar: Mg = 24)
…………….. moles of HCl are present, so …………….. moles of Mg will react with them.
So the mass of magnesium that will react is …………….. g. c ) What volume of 0.05 mol/dm3 potassium manganate(VII) (KMnO4) will be reduced by
25 cm3 of 0.2 mol / dm3 iron(II) sulfate (FeSO4) solution?
Ionic equation: MnO4-(aq) + 5Fe2+ (aq) + 8H+ (aq)
Mn2+ (aq) + 5Fe3+ (aq) + 4H2O (l)
…………….. moles of FeSO4 are used. From the equation, these will react with ……………..
moles of KMnO4. So the volume of potassium manganate(VII) solution reduced
is …………….. cm3.
Electrolysis Electrolyte: Compounds that conduct electricity either in molten state or in aqueous solution are called electrolytes. All ionic compounds are good electrolytes in molten state or in aqueous solution. Covalent compounds are poor electrolytes because they do not contain ions. Electrolysis in molten NaCl When molten NaCl is electrolysed, the positive Na+ migrates to the –ve cathode and is discharged. Reaction at the cathode Na+ + e-­‐ Na (s) Na metal deposits on the cathode. During electrolysis, the negative Cl-­‐ migrates to the +ve anode and is discharged. Reaction at the anode 2Cl-­‐ -­‐ 2e-­‐ Cl2 (g) Cl2 gas is liberated at the anode. Therefore when molten NaCl is electrolysed, it decomposes into Na metal and Cl2 gas. Electrolysis of Dilute H2SO4 (acidified water) H2SO4 2H+ + SO42-­‐ H2O H+ + OH-­‐ Reaction at cathode 2H+ + 2e-­‐ H2 (g) Reaction at anode OH-­‐ is preferentially discharged OH-­‐ -­‐ e-­‐ [OH] [OH] + [OH] H2O + [O] [O] + [O] O2 The overall process is -­‐ -­‐
-­‐
4OH -­‐ 4e 2H2O + O2(g) Electrolysis of dilute sulphuric acid produces oxygen at the anode and hydrogen at the cathode. The Hofmann Voltameter for the Electrolysis of Water Preferential discharge of ions: When two similarly charged ions are in competition in the electrolyte, only one of them is preferred for discharge at the electrode. Which ion is preferentially discharged at the electrode depends on three factors:  Position of the ion in the electro – chemical (E.C) series.  Relative concentration of the ions in the electrolyte.  Nature of the electrode used. Position of the ion in the electro – chemical series The lower the position in the electro – chemical series, greater is its tendency to be discharged at the electrode. Cations Anions K+ Na+ Ca2+ Mg2+ Al3+ Zn2+ Fe2+ Pb2+ H+ Cu2+ Hg+ Ag+ Au3+ Pt+ SO42-­‐ O3-­‐ OH-­‐ Cl-­‐ Br-­‐ I-­‐ K+ Relative concentration of the ions in the electrolyte Higher the concentration of the ion, greater is its probability of being discharged at the electrode. If an ion is in higher concentration, it will be discharged at the electrode even if it is higher than the other in the Electrochemical series. Nature of the electrode used The ionic reaction taking place at the electrode and therefore the product formed depend upon the nature of the electrodes used. Eg. When aqueous CuSO4 is electrolyzed using Pt electrodes, Cu is deposited at the cathode and O2 is liberated at the anode. If the anode is made up of Cu metal, then Cu atoms dissolve to form Cu2+ ions at the anode instead of O2 being liberated. Electrolysis of Copper (II) Sulphate [CuSO4 (aq)] Reaction at cathode 2+
-­‐ Cu (aq) + 2e  Cu (s) Particles of copper are deposited below the cathode. Reaction at the anode 4OH-­‐(aq) – 4e-­‐  2H2O (l) + O2 (g) Oxygen gas is released at the anode. Note -­‐ This process can be used to electroplate articles with copper. The article to be plated is immersed in the electrolyte and made the cathode. The anode is made of pure copper metal and the electrolyte used is CuSO4 soln. The copper ions cannot deposit on the platinum electrodes therefore, they collect as tiny particles below the cathode. As the copper ions get discharged at the cathode the concentration of the Cu2+ ions decrease and the blue colour of the solution fades and finally becomes colourless. If the same electrolysis process is carried out using copper electrodes, the reactions at the electrodes will be different. Reaction at cathode Cu2+ (aq) + 2e-­‐  Cu (s) Particles of copper are deposited on the cathode. Reaction at the anode Cu (s) -­‐ 2e-­‐  Cu2+ (aq) In this case, for every Cu2+ ion discharged at the cathode, an equal number of copper atoms at the anode lose electrons and go into the electrolytic solution as ions. Therefore the blue colour of the electrolytic solution does not fade. Applications of electrolysis 1. Electroplating. 2. Electro-­‐refining of metals. 3. Extraction of metals. Electroplating with silver Reaction at the anode Ag (s) -­‐ e-­‐  Ag+ (aq) Reaction at cathode Ag+ (aq) + e -­‐  Ag (s) During electroplating:  The object to be plated must always be made the cathode.  The pure metal is always made the anode.  The electrolyte must contain the metal ion. Electro-­‐refining of metals Reaction at the anode Cu (s) -­‐ 2e-­‐  Cu2+ (aq) Reaction at cathode Cu2+ (aq) + 2e-­‐  Cu (s) When electrolysis is carried out for a few hours, the anode loses mass as pure copper atoms break down into Cu2+ (aq) ions and go into the electrolytic solution. They migrate to the cathode, gain electrons and deposit at the cathode. Therefore the mass of the cathode increases. At the end all the pure copper metal is deposited at the cathode and the anode will contain only impurities. In this way impure copper is converted into 100% pure metal. Extraction of metals – Extraction of Aluminium from bauxite A mixture of molten alumina (pure Al2O3) and Cryolite (NaAlF6) is taken as electrolyte in the tank and electrolysis is carried out. The temperature is maintained at 9500C. Presence of cryolite lowers the melting point of alumina and helps conduction of electricity. Reaction at cathode Al3+ (l) + 3e-­‐  Al (s) The aluminium metal formed sinks to the bottom and is removed. Reaction at the anode 2O2-­‐ (l) -­‐ 4e-­‐  O2 (g) The oxygen gas liberated combines with the carbon at the anode to form carbon dioxide. This is why a mixture of oxygen and carbon dioxide is found at the anode. Worksheet – 8 Electrolysis 1
a
For A–J below, circle the letter if the bulb will light. Cross it out if the bulb will not light.
b For the substances above where the bulb lights, which will decompose? …………………………………………………………………………………………………… 2
Which row shows the products at the anode and cathode, during electrolysis? Circle its letter. at the anode
A
B
C
D
3
metals
metals
non-metals, including hydrogen
non-metals (except-hydrogen)
at the cathode
non-metals (except-hydrogen)
non-metals, including hydrogen
metals
metals or hydrogen
Which row shows the products from the electrolysis of concentrated sodium chloride solution? Circle its letter. Positive electrode
A
B
C
D
hydrogen
hydrogen
oxygen
chlorine
Negative electrode
oxygen
chlorine
hydrogen
hydrogen
4
Strontium chloride (SrCl2) is melted and electricity is passed through it, using inert electrodes.
Strontium is a reactive metal from Group II of the Periodic Table.
a
Write ionic equations for the reactions at the electrodes:
At the cathode (–):
……………………………………………………………………………
At the anode (+):
…………………………………………………………………...................
b
Name the three products obtained from the electrolysis of concentrated aqueous
strontium chloride.
………………………………………………………………………………………………....
5
The diagrams below are to show apparatus for purifying copper, by electrolysis.
a
To diagram A:
i add a battery and wires to complete the circuit
ii mark + and – on the correct electrodes
b
Complete diagram B to show when electrolysis is almost complete. Mark in the battery, wires, electrodes, electrolyte, and impurities. c
Write the half-­‐equation for the reaction: i at the positive electrode ……………………………………………………………. ii at the negative electrode ……………………………………………………………. d Give one use of copper, that requires it to be very pure: ……………………………….. ………………………………………………………………………………………….. Energy Changes All chemical reactions involve energy changes. Some chemical reactions give out energy in the form of heat, while other reactions absorb heat energy. Chemical reactions that absorb heat are called endothermic reactions. Chemical reactions that give out heat are called exothermic reactions. Energy in the chemical substances are stored in the bonds. Bond – breaking is an endothermic process. Bond – making is an exothermic process. During the process of a chemical reaction, old bonds are broken and new bonds are formed. If the energy required to break old bonds is greater than the energy released when new bonds are formed, the reaction is endothermic. If the energy required to break the old bonds is lesser than the energy released when new bonds are formed, then the reaction is exothermic. Example: 2H2 + O2 2H2O (exothermic) Burning of Hydrogen is an exothermic process because the energy required to break H – H bond and the O = O bonds is less than the energy released when the two O – H bonds are formed. N2 + O2 2NO (endothermic) In this reaction, the energy required to break the bonds and bonds is much greater than the energy released when two are formed. Energy is measured in Joules or Kilojoules. During a chemical process, the energy change is stated as joules/kilojoules per mole. Fuels Any substance, which on burning produces large amounts of energy, can be used as a fuel. The most commonly used fuel in modern times are the hydrocarbons. Hydrocarbons are made of Carbon and Hydrogen only. Example: Methane – CH4 (g), Ethane – C2H6 (g), Propane – C3H8 (g) etc. the other higher hydrocarbons are present in fuels like petrol, kerosene, diesel etc. Hydrogen is a very good fuel as it burns giving out large amounts of heat energy. It is also a non – polluting fuel as the product of burning Hydrogen is water vapour. The only major problem in using hydrogen is the explosive nature of burning. Alternate Energy Sources 1. Wind energy 2. Tidal energy 3. Solar energy 4. Biogas Out of these, solar energy and biogas have practical use. Solar energy is the best alternate source of energy, as it produces no pollution. The only problem with solar energy is the high manufacturing cost. The other best alternate fuel is biogas, also called natural gas. It is a mixture of hydrocarbons but mostly contains methane. It is formed by the anaerobic degradation of organic matter. All organic waste, under the right anaerobic condition, can be used to produce biogas or natural gas. The only problem with biogas is it is poisonous as it contains high levels of methane. Nuclear energy Nuclear energy is obtained by nuclear fission or by nuclear fusion. Nuclear fission is easier to carry out in nuclear reactors and therefore, is widely used as a source of energy. Nuclear fusion is very difficult to initiate, as it requires very temperatures, which is impossible to duplicate on earth. Nuclear fusion occurs in the stars, which radiate as heat and light. Nuclear fission When a slow neutron is fired into the nucleus of a Uranium – 235 atom, the nucleus splits into a Barium nucleus and a Krypton nucleus releasing large amounts of energy. 235
137
+ n1 + 36Kr84 + 3n1 + energy 92U
56Ba
‘The breaking of a heavier atom into two smaller atoms with the liberation of large amounts of energy is called nuclear fission.’ When the Uranium nucleus splits, it releases neutrons along with energy. These neutrons further split other Uranium atoms forming a chain reaction. The chain reaction is so fast that the whole mass of Uranium explodes releasing large quantities of energy in the form of heat and light. If exploded without control, it is a nuclear bomb. However, if the chain reaction is slowed down in a nuclear reactor and controlled, the energy released can be used to produce electricity. Small masses of Uranium can give large quantities of energy. That is why nuclear energy is much sought after. Dangers of using nuclear fuel  Large amounts of radiation are produced which if leaked can cause tremendous harm to life around the reactor.  The spent fuel left behind after nuclear fission cannot be easily disposed off because it is radioactive waste and can harm the environment.  Nuclear reactors are difficult to construct and maintain as it involves sophisticated advanced engineering. Nuclear fusion Nuclear fusion is the joining of two light nuclei to form a heavier nucleus with the release of large amounts of energy. 2
3
4
1H + 1H 2He + energy Bringing two light nuclei together to form a large nucleus and energy is known as fusion. Nuclear fusion can be a large source of energy, but it cannot be artificially carried out as large amounts of energy is initially required to bring the two light nuclei together to form a large, heavy nucleus. Worksheet – 9 Energy changes 1 The apparatus on the right is used to measure temperature changes during
reactions in solution. Dilute hydrochloric acid is placed in the polystyrene cup
and its temperature recorded. Some of solid P is added.
The mixture is continuously stirred and the temperature is checked regularly.
When there is no further change in temperature, the final temperature is recorded.
thermometer
polystyrene
cup and lid
reaction
mixture
The experiment is repeated with solid Q.
This table shows the results:
Solid
Initial
Final
Temperature temperature/ °C
/ °C
P
20
33
Q
20
12
Temperature
change / °C
a) Complete the table to show the temperature change.
(Write + or – before each value.)
b) i Why is it important to stir continuously during the experiment?
……………
…………………………………………………………………………….
ii If the solution is put directly into the glass beaker, without the polystyrene cup,
how will the results be affected?
..…………………………………………………………………………………………
...………………………………………………………………………………………..
c) Complete these energy diagrams for the two reactions, by marking in the lines for
the products, and up or down arrows to show the energy change:
2 You can compare the energy given out by different fuels, using
the apparatus on the right.
a) What change will you observe, which confirms that the
reaction is exothermic?
……………………………………………………………
……………………………………………………………
b) What precautions will you take, to make sure you are comparing the fuels in a
fair test?
………………………………………………………………………………….
………………………………………………………………………………….
c) Suggest one change to the apparatus, that will improve the accuracy of the comparison. …………………………………………………………………………………………
d) Complete and balance this equation for the combustion of ethanol, when it is used as the fuel: C2H5OH + O2
Rates of Reactions Rate of a reaction -­‐ is change in concentration of the reactants or products with time. If one of the products is a gas, then the volume of gas collected and the total time taken is measured to calculate the rate. The gas liberated during the reaction is measured using a gas syringe. Factors that affect the rate of a reaction  Concentration of the reactants  Temperature  Surface Area (Size of particles)  Catalyst  Pressure (for gaseous reactants)  Light Intensity (for photochemical reactions) Concentration: Greater the concentration of the reacting substance, greater is the number of particles present. Therefore, there is more collision between the reactant particles, increasing the rate of the reaction. Temperature: Higher the temperature, greater the K.E of the particles. Greater the K.E, greater will be the collisions between the particles therefore faster will be the reaction. Surface Area: Smaller the particles, greater is their surface area. Greater the surface are, faster is the chemical reaction. This is because there is greater interaction between the reactant molecules. Catalyst: Presence of catalyst lowers the energy required to break the old bonds. Therefore, presence of catalyst increases the rate of a reaction. Pressure: Increase in pressure decreases the volume of gases. Therefore, closer will be the reactant molecules and faster will be the reaction. Intensity of light: In photochemical reactions, theintensity of light increases the light energy available for the reaction. Therefore, greater the intensity, faster is the rate of photochemical reaction. Measuring Rate of a Reaction in the Laboratory Effect of temperature: -­‐ Na2S2O3 (aq) + 2HCl (aq) Sodium Thiosulphate 2NaCl (aq) + H2O (l) + S (s) + SO2 (g) In the above reaction, sulphur particles are precipitated. The reaction produces solid particles of Sulphur which precipitate. Because of this precipitation the cross on the paper disappears after some time. The time taken for the cross to disappear measures the rate of the reaction. Effect of temperature on rate of reaction: -­‐ The above experiment is carried out at different temperatures. And the time taken for the cross to disappear is measured. It is found that at higher temperatures the cross disappears faster. This proves that rate of a reaction increases with temperature. Effect of concentration on rate of a reaction: -­‐ increase in the concentration of the reactant molecules increases the rate of reaction. This is because increase in concentration increases the number of molecules. This increases the number of collisions between the reactant molecules, thereby increasing the rate of the reaction. Effect of catalyst: -­‐ A catalyst is a substance which increases the rate of a reaction without undergoing any permanent damage. Addition of a catalyst to the reactants increases the rate of reaction. This is because the catalyst lowers the ‘energy of activation’. The energy of activation is defined as the minimum energy required to make the reaction happen. The presence of a catalyst increases the rate of the reaction by lowering the energy of activation for the process. The reactant molecules form temporary bonds with the catalyst and therefore their bonds are weakened. As a result they require less energy to break. Example: Fe in Haber’s process (manufacture of NH3), V2O5 in contact process (manufacture of H2SO4), MnO2 in preparation of O2 from H2O2 Effect of pressure: -­‐ for gaseous reactants, increase in pressure increases the rate of reaction. This is because increase in pressure decreases the volume of the gas. Under decreased volume the reactant molecules collide faster and therefore rate of reaction increases. Photochemical reactions: -­‐ these are chemical reactions which take place only in the presence of light. Eg. Photosynthesis, Photography. Photography is the process of redox reaction using AgCl2 or AgBr2. 2AgBr 2Ag + Br2 -­‐
2Br Br2 + 2e – (oxidation) Ag + + e -­‐ Ag(s) (reduction) The redox process depends on the amount of light present. That is why it is called photochemical reaction. Greater the intensity of light, faster will be the photochemical reaction. Worksheet – 10 Rates of reaction 1
Nitrogen gas is insoluble in water. It is produced in the reaction between warm solutions
of ammonium chloride and sodium nitrite:
NH Cl (aq) + NaNO (aq)
NaCl (aq) + 2H O (l ) + N ( g)
The rate of the reaction can be followed by measuring the volume of gas given off, over time.
4
2
2
2
a)
The gas can be collected over water. Complete this diagram to show the apparatus needed:
b)
This table shows the results when the reaction was carried out at 70 °C.
Time / min
0
1
2
3
4
5
6
7
8
9
10
Volume / cm3
0
20
32
43
45
54
57
58.5
59.5
60
60
`
Plot a graph volume against time, on a piece of graph paper. Join the points with a smooth
curve that fits the points best. Then attach your graph to this worksheet.
c)
i Using your graph to help you, describe how the rate of the reaction varied with time:
.…………………………………………………………………………………………………
.…………………………………………………………………………………………………..
ii Why did the rate vary?
.....................................................................................................................................................
..…………………………………………………………………………………………………
d)
i One of the results was anomalous – it did not fit the pattern. Circle that point on your graph.
ii Suggest a reason for the anomalous result:
………………………………………………………………………………………………..
.…………………………………………………………………………………………………
……………………………………………………………………………………………………
e
The reaction was repeated at 80 °C. Draw a rough sketch of the graph you would expect to
obtain, on the same axes as your graph in b.
2
An experiment was carried out to investigate the rate of reaction
between dilute hydrochloric acid and an excess of calcium
carbonate, in the form of marble chips:
CaCO3 (s) + 2HCl (aq)
CaC2 (aq) + H2O (l ) + CO2 (g)
a What will you observe in the flask, when the acid is added
to the marble? ...................................................................................
……………………………………………………………………….
b What is the purpose of the cotton wool?
……………………………………………………………………………………………………
c Why is there a loss of mass as the reaction proceeds?
.........................................................................................................................................................
……………………………………………………………………………………………………
3
The experiment in 2 was repeated, keeping everything the
same except the concentration of the acid. Two different
concentrations were used, A and B. Look at the results in
this table.
concentration
of the acid
loss of mass in
first minute
A
0.5g
B
1g
a
Which concentration was higher, A or B? ..........................
b
Explain why one reaction was faster, in terms of collisions between reacting particles:
……………………………………………………………………………………………………………
…..
……………………………………………………………………………………………………………
…..
4
The experiment in 2 was repeated, changing
only the initial temperature of the acid.
Two temperatures were used, C and D.
The results are shown in this table.
initial temperature
of the acid (°C)
loss of mass in
first minute
C
0.5 g
D
2g
a
Which was higher, C or D? ...............................
b
Explain why one reaction was faster, in terms of collisions between reacting particles:
…………………………………………………………………………………………………
…………………………………………………………………………………………………
Chemical Equilibrium N2 + 3H2 2NH3 H = -­‐92kJmol-­‐1 In the above chemical reaction, there is a dynamic equilibrium. The rate of the forward reaction is equal to the rate of the backward reaction. Therefore, left to itself there is a balance between the forward and backward reactions. We say the equilibrium is at the centre. If any of the factors on which the equilibrium depends is altered, the equilibrium shifts either to the right or to the left neutralizing the effect of the change. In the manufacture of NH3 by Haber’s process, high pressure and low temperature favours the formation of NH3 as the equilibrium shifts to the right. i.
N2 + 3H2 2NH3 ∆H = -­‐92 kJmol-­‐1 Low temperature (350˚C -­‐ 450˚C) High pressure (200 Atm) ii.
2SO2 + O2 2SO3 ∆H = -­‐197 kJmol-­‐1 Low temperature (450˚C) High pressure (1 -­‐ 2 Atm) iii.
N2 + O2 2NO ∆H = +90 kJmol-­‐1 High temperature (3000˚C) No effect of pressure For endothermic reactions: Increase in temperature shifts the equilibrium to the right (products side). For exothermic reactions: Increase in temperature shifts the equilibrium to the left (reactants side). Worksheet – 11 Chemical Equilibrium
1 Sulphuric acid is made by the Contact Process. 2SO2(g) + O2(g) -­‐1
2SO3(g) ΔH = -­‐ 197 KJ mol
(i) What are the reaction conditions for the Contact Process? .....................................................................................................................................................
............... .....................................................................................................................................................
............... (ii) Would the yield of sulphur trioxide increase, decrease or stay the same when the temperature is increased? Explain your answer. .....................................................................................................................................................
................ .....................................................................................................................................................
................ .....................................................................................................................................................
................ ……................................................................................................................................................
............ 2. Ammonia is made on a large scale by the Haber’s process. N2(g) + 3H2(g) 2NH3(g) ΔH = -­‐ 92 KJ mol-­‐1 a) State two characteristics of a reversible reaction at equilibrium. .....................................................................................................................................................
........... .....................................................................................................................................................
........... b) How can the amount of ammonia in the reaction mixture be increased? .....................................................................................................................................................
................ ……................................................................................................................................................
............ Oxidation and Reduction Oxidation and Reduction take place simultaneously. In terms of electron transfer – Oxidation is loss of electrons Remember -­‐ OIL RIG Reduction is gain of electrons Metals have a tendency to lose electrons and therefore undergo oxidation. Non – metals have a tendency to gain electrons and therefore undergo reduction. Metals donate electrons and cause reduction to take place in other substances. Therefore, they are reducing agents or reductants. Non – metals gain electrons from other substances and cause oxidation to take place. Therefore, they are oxidising agents or oxidants. Oxidising agents (Oxidants) Other than Non – metals 1. Hydrogen Peroxide – H2O2 2. Potassium Manganate (VIII) – KMnO4 3. Potassium Dichromate – K2Cr2O7 These substances release nascent oxygen [O] Reducing agents (Reductants) Other than Metals 1. Hydrogen – H2 2. Carbon – O2 3. Carbon Monoxide – CO 4. Sulphur Dioxide – SO2 5. Ammonia – NH3 Example: -­‐ 1) CuO + H2 Cu + H2O In this reaction, H is being oxidised to H2O. The oxidising agent is Cu. CuO is being reduced to Cu. The reducing agent is H2. 2) CH4 + 2O2 CO2 + 2H2O CH4 is oxidised to CO2 and H2O. O2 is the oxidising agent. O2 is being reduced to H2O. Conversion of FeCl2 to FeCl3 FeCl2 + ½ Cl2 FeCl
3 Iron (III) Chloride Iron (II) Chloride When Cl2 is passed through aqueous FeCl2, it oxidises FCl2 to Fe Cl3 Fe2+ -­‐ e-­‐ Fe3+ (oxidation) ½ Cl2 + e-­‐ Cl-­‐ (reduction) Colour change: FeCl2 + ½ Cl2 FeCl3 Pale green Orange yellow Potassium Manganate (VIII) – purple/pink When reduced, it becomes pale pink Potassium Dichromate – orange When reduced, it becomes green Worksheet – 12 Oxidation- Reduction 1
When hydrogen is passed over copper(II) oxide, this reaction takes place:
CuO (s) + H2 (g)
Cu (s) + H2O (l)
a
It is a redox reaction, because ………………………………………………………………...
……………………………………………………………………………………………….......
b
The reducing agent in this reaction is …………………………………………………………
……………………………………………………………………………………………….......
2
Consider the following reaction : Fe2O3 (s) + 3CO ( g)
2Fe (l) + 3CO2( g)
a
The word equation for the reaction is:
……………………………………………………………………………………………….......
.
b
It is a redox reaction, because …………………………………………………………………
…………………………………………………………………………………………………
…
c
The reducing agent in this reaction is ………………………………………………………....
3
A coil of copper wire (Cu) is placed in a colourless solution of
silver nitrate (AgNO3). The solution changes colour.
a
What colour does the solution go, and why?
…………………………………………………………
…………………………………………………………
Write an ionic equation for the reaction that takes place.
b
……………………………………………………………………………………………………………
c The copper is said to be oxidised during this reaction. Explain why.
……………………………………………………………………………………………………………
4
Explain why this is not a redox reaction: CuO (s) + H2SO4 (aq)
CuSO4(aq) + H2O (l)
……………………………………………………………………………………………………………
5 Group
I
II
III
IV
V
VI
VII
O
Element
sodium
magnesium
aluminium
silicon
phosphorus
sulfur
chlorine
argon
Typical
Compound
NaCl
MgO
AlCl3
SiO2
PH3
H 2S
HCl
none
Oxidation
state of
element in it
+I
+II
+III
+IV
–III
–II
–I
–
a
The table above shows a period of the Periodic Table. Which period? ....................
b
The oxidation state tells you how many electrons an atom has ...................., ...................., or
...................., in forming a compound.
c
i Why are the oxidation states of the elements in the compounds negative, after Group IV?
………………………………………………………………………………………………
ii Why do they decrease, after Group IV?
………………………………………………………………………………………………
d
Using oxidation states, write the formulae of the compounds formed between:
………….
i
sodium and sulfur ………….
ii
silicon and chlorine ………….
iii
aluminium and sulfur ………..
iv
magnesium and phosphorus
Acids, Bases and Salts An acid is a substance that produces H+ ions in aqueous solution. Examples: HCl, HNO3, H2SO4, H2CO3, H3PO4 A base is a substance that reacts with acids to form salt and water only. Examples: NaOH, KOH, Ca(OH)2, NH4OH A salt is a substance formed by the neutralisation of an acid and a base. Examples: NaCl, KCl, CaCO3, NH4Cl, NaNO3 Neutralisation reactions HCl + NaOH NaCl + H2O Acid Base Salt Water HCl + NH3 NH4Cl + H2O Acid Base Salt Water H2CO3 + CaO CaCO3 + H2O Acid Base Salt Water
During neutralisation the H+ ions are neutralise by the OH-­‐ ions. H+ + OH-­‐ H2O A salt consists of a positive ion called the cation and a negative ion called the anion. The cation comes from the base and the anion comes from the acid. The pH Scale It is a scale ranging from 1 to 14 that is used to measure the acidity/ basicity of the substance. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Acid Alkali Neutral The lower the pH value stronger is the acid. The higher the pH value stronger is the base. When the concentration of H+ ions and the OH-­‐ ions are equal the pH is 7. The pH of a solution can be measured by using a pH meter. Preparation of salts Salts can be prepared by two ways: 1. By neuralisaton. 2. By precipitation. Preparation of salts by neutralisation Soluble salts are prepared by neutralisation. Copper Sulphate CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l) Sodium Chloride HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) Method: 1. Take the acid in a beaker. 2. Add the base to it slowly and heat it. 3. When the reaction is complete allow it to cool. 4. Filter the mixture to remove the impurities. 5. The filtrate obtained is the salt solution. 6. Take the filtrate in a crucible and heat it to evaporate the water 7. Allow it to cool. The crystals of salt are formed. Preparation of salts by precipitation [mixing solutions of two soluble salts] Insoluble salts are prepared by precipitation. Exam,ple 1-­‐ Lead Chloride Pb(NO3)2(aq) + 2HCl(aq) PbCl2(s) + HNO3(aq) Example 2-­‐ Barium Sulphate Ba(NO3)2(aq) + H2SO4(aq) BaSO4(s) + 2HNO3(aq) Insoluble salts can be prepared by mixing two salt solutions containing the required cation and anion. In the above examples lead chloride is prepared by mixing aqueous solutions of lead nitrate which provides the cation Pb2+ and aqueous hydrochloric acid which provides the anion Cl -­‐. Similarly in the next example Ba2+ cation comes from barium nitrate solution while the SO42-­‐ anion comes from sulphuric acid. In this way other insoluble salts can be prepared. Method: 1. Choose soluble salts containing the required cation and anion. 2. Make solutions of the two salts in two separate beakers. 3. Mix the two solutions in a third beaker 4. The insoluble salt will precipitate. 5. Filter the mixture. 6. The insoluble salt remains in the filter paper. 7. Wash the salt prepared with distilled water. 8. Allow it to dry. Indicators Indicators are compounds, which show colour changes in acidic and alkaline medium Examples: Litmus solution, methyl orange, phenolphthalein, universal indicator etc. Indictor Colour Acid Medium Alkaline Medium Litmus Methyl Orange Phenolphthalein Universal Indicator Violet Orange Colourless Green Red Red Colourless Red Blue Yellow Pink Blue (or) Violet TYPES OF OXIDES Oxides can be acidic, basic or neutral. Metal oxides are all basic. Some metal oxides are amphoteric. Non-­‐metal oxides are all Acidic or Neutral. All metal oxides are basic in nature. Some metal oxides like Al2O3, PbO and ZnO are amphoteric oxides. (They can react with both acids as well as strong alkalis to form salt and water). Example: CaO + H2SO4 CaSO4 + H2O (Base) (Acid) (Salt) (Water) Al2O3 + 6HCl 2AlCl3 + 3H2O (Base) (Acid) (Salt) (Water) The same Al2O3 can react with strong alkalis to form a salt (Sodium Aluminate) and water. Al2O3 + 2NaOH 2NaAlO2 + H2O (Acid) (Base) (Salt) (Water) Similarly, ZnO and PbO show amphoteric nature. ZnO + H2SO4 ZnSO4 + H2O (Base) ZnO + 2NaOH Na2ZnO2 + H2O (Acid) (Base) (Sodium Zincate) PbO + H2SO4 PbSO4 + H2O PbO + 2NaOH Na2PbP2 + H2O (Base) (Acid) (Base) (Sodium Plumbate) Non – metal Oxides like CO2, SO2, SO3, and NO2 are acidic in nature. However, CO, NO and H2O are neutral in nature. NH3 is the only alkaline gas. Acid Rain When gases like SO2, SO3 and NO2 are released into the atmosphere, they dissolve in rainwater and come down as acid rain. SO2 + H2O H2SO3 (Sulphurous acid) SO3 + H2O H2SO4 (Sulphuric acid) 4NO2 + 2H2O + O2 4HNO3 (Nitric acid) Worksheet – 13 Acids and Bases 1 To prepare the salt potassium chloride, 25 cm3 of potassium hydroxide were first
titrated against a solution of acid, with phenolphthalein as indicator.
Phenolphthalein is colourless in acid solution, and pink in alkaline solution.
a The drawings below show the pieces of apparatus used.
i Name each piece.
ii
Say what was placed in each, during the titration.
i ………………………………. i ………………..................
i ………………………............
ii ……………………………… ii …………………………….. ii ……………………………..
b Describe the colour change that took place, showing that neutralisation was complete.
……………………………………………………………………………………………………………
c The first burette on the right shows the initial reading for the titration.
The second shows the final reading. Use them to complete this table.
initial reading /
cm3
final reading /
cm3
volume used /
cm3
d To prepare the salt, the titration was then repeated. But there was one important
change. State the volumes of acid and alkali that were used, and the change that was made.
……………………………………………………………………………………………………………
e
The final stage was to evaporate water from the mixture obtained in d. Why was this done?
……………………………………………………………………………………………………………
……………………………………………………………………………………………………………
f
Write the word equation for the neutralization reaction that produced the salt.
…………………………………………………………………………………………………………….
2
Below is a list of twelve salts, in alphabetical order. There are four insoluble salts, and four pairs
of soluble salts from which these insoluble salts can be made.
barium chloride, barium sulfate, calcium carbonate, calcium nitrate, lead iodide, lead nitrate
potassium iodide, potassium sulfate, silver nitrate, silver chloride, sodium carbonate, sodium
chloride
a
Complete this table using salts from the list.
This insoluble salt …
can be made using this
…
and this …
1
2
3
4
b
The method for making an insoluble salt from two soluble salts is called ……………………
c
i
In the table below, write ionic equations for the reactions that produce the four insoluble
salts. Include the state symbols (but not the spectator ions).
ii Then give the formulae for the two spectator ions present at each reaction.
Ionic equation for the reaction
3
Spectator ions
1
and
2
and
3
and
4
and
These are all oxides: dinitrogen oxide calcium oxide phosphorus oxide zinc oxide
a Write their names in the correct places in this table, and add their chemical formulae.
acidic oxide
basic oxide
neutral oxide
amphoteric oxide
acidic oxide
basic oxide
neutral oxide
amphoteric oxide
name
formula
b i
Which of the four oxides will react with sodium hydroxide? ……….......................................
ii Which of them will react with hydrochloric acid? .....................................................................
iii Explain the term amphoteric …………………………………………………………………..
………………………………………………………………………………………………….
METALS GENERAL PROPERTIES  Most metals are hard and have lustre  They have high melting and boiling points  They are malleable and ductile  Metals generally have high density  They are good conductors of heat and electricity  Most metals have high tensile strength Metal Reactivity Series K+ (most reactive) Na+ Ca2+ Mg2+ Al3+ Zn2+ Fe2+ Pb2+ Cu2+ Hg+ Ag+ Au3+ Pt+ (least reactive) The metal reactivity series K+ is an arrangement of metals in order of their reactivity beginning with the most reactive and ending with the least reactive metal. Reactive metals like K, Na, Ca, Mg and Al are all extracted by electrolysis from their compounds. These metals cannot be extracted from their oxides by heating with coke. Less reactive metals like Zn, Fe, Pb and Cu are extracted from their ores by chemical methods. These metals can be obtained from their oxides by heating with coke. Metals like Au and Pt exist in their native states and are obtained from rocks by separation techniques. Main steps involved in the extraction of metals from their ores •
Concentration of the metal ore •
Roasting/Calcination of the concentrated ore •
Reduction of the metal oxide •
Refining of the metal EXTRACTION OF ALUMINIUM FROM BAUXITE (Al2O3) 1. Bauxite is concentrated by chemical process to get pure Alumina (Al2O3) 2. Alumina is mixed with Cryolite (Na3AlF6) and electrolysed in molten state at 950ºC ( - )
Graphite Carbon rods Reaction at cathode Al3+ (l) + 3e-­‐  Al (s) The aluminium metal formed sinks to the bottom and is removed. Reaction at the anode 2O2-­‐ (l) -­‐ 4e-­‐  O2 (g) (+)
Molten
Al2O3 and
Cryolite
The oxygen gas liberated combines with the carbon at the anode to form carbon dioxide. This is why a mixture of oxygen and carbon dioxide is found at the anode. To prevent the burning of the Carbon rods, a layer of coke is sprinkled on the surface of the electrolyte. 3. Purification – This step is not required in this process as the Aluminium obtained is over 99% pure. SACRIFICIAL PROTECTION When a more reactive metal is in contact with another metal, the more reactive metal corrodes protecting the other metal. This is known as Sacrificial Protection. Example: When Zinc metal is in contact with Iron, it corrodes protecting the Iron from rusting. This is the reason why strips of Zinc are riveted to the bottoms of ships made of steel. The Zinc metal strips slowly corrode protecting the steel body of the ship. This is the reason why iron sheets are galvanised (coated with Zinc). Worksheet – 14 Metals 1
The first column in the table below lists some general properties of metals.
a
Complete the second and third columns of the table. Write neatly!
General property of metals
Correct name for this
property
One use that depends
on this property
can be drawn into wires
can be bent into shape
reflect light
make a ringing sound when struck
allow electricity to pass through
heavy for their volume
their oxides react with acids
transfer heat well
b i Which property above is a chemical property? ..................................................................................................................................................
ii Some chemical properties apply only to metals high in the reactivity series.
Give one example:
…………………………………………………………………………………………...
iii When metals react, they form ions with a ………………………………. charge.
2
Going down Group II in the Periodic Table, you will find magnesium, calcium and strontium,
in that order. Look at these observations:
Observation 1: Magnesium reacts very slowly with cold water, but more vigorously with
steam.
Observation 2: Calcium reacts briskly with cold water, with a lot of fizzing.
a
Those two observations show that reactivity ……………………………. down Group II.
bi
Predict how strontium will react with cold water.
……………………………………………………………………………………………
ii
Which gas is released, in the reaction of those metals with water?
...........................................................................................................................................
c
One Group II element reacts with neither water nor steam. Which one?
...........................................................................................................................................
d
Magnesium reacts with oxygen to form magnesium oxide (MgO). What forms when
strontium reacts with oxygen? Give its name and formula.
............................................................................................................................................
e
When magnesium oxide is heated with carbon, no reaction occurs. So carbon is …
………………………magnesium in the reactivity series.
ORGANIC CHEMISTRY The carbon atom can form four covalent bonds with other atoms. It also has the ability to repeatedly bond with other carbon atoms. This ability is known as catenation. Homologous Series It is a series of compounds that have a general formula and a common functional group, with each member in the group differing from the other by a -­‐CH2 group. Example: Alkanes, alkenes, alkynes, alcohols etc. Compounds that form a homologous series have similar chemical properties and gradation in physical properties. Alkanes (CnH2n+2) Alkanes are the simplest family of hydrocarbons -­‐ compounds containing carbon and hydrogen only. They only contain carbon-­‐hydrogen bonds and carbon-­‐carbon single bonds. The first six are: Methane CH4 Ethane C2H6 Propane C3H8 Butane C4H10 Pentane C5H12 Hexane C6H14 Physical Properties Boiling Points -­‐ The boiling points of the alkanes increase with molecular size. Long chained molecules have higher boiling points. Fractional distillation of petroleum (crude oil) -­‐ Petroleum consists of a mixture of many hydrocarbons. When petroleum is fractionally distilled, it separates out into many fractions. These fractions are useful substances as fuel and raw material for chemical industry. Petroleum is vaporized and let into the fractionating column as shown above in the diagram. The vapours cool as they rise up through the column. The different components of petroleum condense at different levels in the fractionating column. In this way, many components from the mixture can be separated simultaneously. The major components obtained from petroleum are -­‐ 1. Petroleum gas (LPG) 2. Petrol (Gasoline) 3. Naphtha 4. Kerosene (Paraffin) 5. Diesel 6. Lubricating oil (Engine oil) 7. Fuel oil 8. Bitumen Solubility -­‐ Alkanes are insoluble in water, but dissolve in organic solvents. The liquid alkanes are good solvents for many other covalent compounds. Chemical Properties Alkanes contain strong carbon-­‐carbon single bonds and strong carbon-­‐hydrogen bonds. They are relatively inert and react with difficulty. However they are good fuel. Combustion -­‐ Complete combustion of any hydrocarbon produces carbon dioxide and water. Halogenation (photochemical reaction) In the presence of ultra-­‐violet light /sunlight alkanes undergo substitution reaction. There is no reaction in the dark. The reactions with bromine are similar. Cracking -­‐ When .long chained hydrocarbons are heated in the presence of a catalyst, they break down into smaller chain hydrocarbons along with the formation of ethene. This is known as Catalytic Cracking. Decane
Hexane
Ethene Alkenes (CnH2n) Alkenes are a family of hydrocarbons (compounds containing carbon and hydrogen only) containing a carbon-­‐carbon double bond. The first four are: Ethene C2H4 Propene C3H6 Butane C4H8 pentene C5H10 Chemical Properties Alkenes are very reactive and undergo addition reactions. Hydrogenation -­‐ Ethene reacts with hydrogen in the presence of a finely divided nickel catalyst at a temperature of about 150°C to form ethane. Addition of bromine -­‐ Bromine adds on to ethene to give 1,2-­‐dibromoethane. The reddish-­‐brown colour of bromine is decolourised as it reacts with the alkene. Using bromine water as a test for alkenes -­‐ If you bubble a gaseous alkene through bromine water, the solution becomes colourless. Alkenes decolourise bromine water. This is a test for unsaturation. Addition of hydrogen halides -­‐ When ethene and hydrogen chloride react, you get chloroethane: Addition of steam to alkene -­‐ Ethanol is manufactured by reacting ethene with steam. Similarly you will get propan-­‐2-­‐ol when propene reacts with steam. Hydrogenation of fats -­‐ Vegetable oils (fats) are unsaturated fats. They have two or more double bonds in their molecular structure. When hydrogen gas is passed through these oils in the presence of Nickel catalyst, the oils get hydrogenated and form margarine (saturated fats). The hydrogen atoms add on across the double bonds converting them into saturated fats. Isomerism -­‐ Isomers are molecules that have the same molecular formula, but have a different structure or arrangement of the atoms in space. For example, there are two isomers of butane, C4H10. Pentane, C5H12, has three chain isomers. In alcohols such as C4H9OH there are two isomers. Alcohols (CnH2n+1OH) Alcohols are compounds similar to hydrocarbons in which a hydrogen atom is substituted by a – OH group. Example: Methanol CH3 – OH Ethanol C2H5 -­‐ OH Propanol C3H7 – OH Manufacture of Ethanol -­‐ In industries, ethanol can be manufactured by reacting ethene and steam. + Ethene Steam Ethanol Fermentation -­‐ The conversions of sugars into alcohol by the action of yeast cells in the absence of oxygen is known as fermentation. Yeast cells anaerobically respire and derive energy. During the process, sugars are converted to alcohol. (Glucose) (Ethanol) Oxidation of alcohols -­‐ When alcohols are oxidised, they form carboxylic acids. Example – Ethanol gets oxidised to Ethanoic acid. When wine is exposed to air, it undergoes slow oxidation to form vinegar. This is chemically the oxidation of Ethanol to Ethanoic acid. Formation of Esters -­‐ Alcohols react with carboxylic acids in presence of H+ ions slowly to form esters. Esters are a group of compounds, which have a fruity smell. When Ethanol is reacted with Ethanoic acid in presence of Hydrogen ions, Ethyl Ethanoate and Water are formed. (Ethanol) + (Ethanoic acid) → (Ethyl Ethanoate) 1
Paraffin is one of the fractions distilled from petroleum. What is it used for? Tick the box.
2
A as bottled gas, for cooking
B as aircraft fuel
C as fuel for power stations
D as a lubricant
Which of these compounds has the formula C2H4? Tick its box.
A ethane
3
Worksheet – 15 Organic chemistry B ethane
C ethanol
D ethanoic acid
Long-chain alkanes are often cracked to produce more useful products.
a Give two reasons why long-chain alkanes are not very useful.
..........................................................................................................................................................
.............................................................................................................................................................
b When the liquid alkane decane was cracked, a gas formed. It turned bromine water
colourless.
i
What can you deduce about this gas?
............................................................................................
c
ii
The reaction between the gas and bromine is called an ................................... reaction.
i
This diagram shows part of the apparatus for cracking decane in the lab. Complete it.
ii
What is the aluminium oxide for?
...............................................................................................
iii
Mark on the diagram where the test-tube should be heated.
iv
The moment heating is stopped, the apparatus is removed from the water. Why?
....................................................................................................................................................
d
Complete this equation for the cracking of decane, and name the gas that formed:
C10H22
decane
e
C8H18 + ………………………
octane
Complete this table comparing decane with the gas that formed, for the reaction in d:
Compound
Organic family
decane (a liquid)
…………………… (a gas)
Is its boiling point above, or below,
room temperature?
Does it contain double bonds?
Is it saturated or unsaturated?
Will it react with bromine water?
Will it polymerize?
4
Ethanol is a member of the alcohol family.
a
i
What is the functional group of the alcohol family?
..................................................................................................
ii
Write the formula for ethanol
...................................................................................................
b
Ethanol can be made by the fermentation
of sugar, using the apparatus on the right.
The fermentation takes place in the flask.
i
What is put in the flask?
............................................................................................................
ii
Which of these temperatures is best for the reaction? Circle your choice.
0 °C
iii
10 °C
25 °C
75 °C
What would you observe in the test-tube?
..............................................................................
iv
Complete the equation for the fermentation:
. C6H12O6
v
How would you separate the ethanol from the mixture in the flask?
...................................................................................................................................................
c Ethanol is also made from ethene, in an addition reaction.
i
Give the balanced symbol equation for the reaction.
..................................................................................................................................................
ii
Name the catalyst used to speed up the reaction.
..................................................................................................................................................
POLYMERISATION Polymerisation is the process by which small molecules called monomers are linked to form large macromolecules called polymers. There are two types of polymerisation possible –  Addition Polymerisation  Condensation Polymerisation ADDITION POLYMERISATION -­‐ This type of polymerisation takes place by the breaking of one of the double/triple bonds in a molecule to form long chained polymers. Formation of polyethene
Formation of polychloroethene Formation of polytetrafluoroethene In addition polymerization only one type of monomer is used. The breaking of one covalent bond releases a pair of electrons which can form covalent bonds with other monomers to form long chains. Any molecule which has a double or triple bond between the carbon atoms can be used as monomers to produce addition polymer. CONDENSATION POLYMERISATION In this type of polymerization two different monomers are used to form long chained polymer by the elimination of a water molecule. This is why this type of polymerization is called condensation polymerization. There are two main groups of condensation polymers. 1. Polyamides. Example – Nylon When a dioic acid and a diamine are polymerized they form a polyamide. The linkage is called amide linkage. Nylon is a synthetic polyamide while proteins are natural polyamides. 2. Polyesters. Example – Terylene When a diol and a dioic acid are polymerized a polyester is formed. The linkage is called ester linkage. Terylene is a synthetic while fats are natural polyesters. NATURAL POLYMERS Proteins, fats and carbohydrates are all natural polymers found in living systems. These are the three classes of compounds which form the main constituents of food. Proteins Proteins are formed by the polymerization of amino acids. These amino acids are joined together by amide linkage. Digestion of protein is a hydrolysis reaction where water molecules used in the presence of enzymes break the amide linkage. During this process the protein chain is broken down into individual amino acids. One hydrogen atom from the water molecule bonds with the nitrogen atom, while the OH from the water bonds with the carbon atom of the amino acid. Carbohydrates Carbohydrates are polymers of simple sugars like glucose(C6H12O6). The glucose molecules are linked together by the glycosidic linkage to form different types of carbohydrates. HO -­‐ -­‐OH + HO -­‐ -­‐OH Digestion of carbohydrates is a hydrolysis process in which water molecules in the presence of enzymes break the glycosidic linkage to form glucose molecules Fats Fats are polymers formed by the polymerization of one glycerol molecule with three fatty acids. It has ester linkage just like in teryline. Digestion of fat is hydrolysis where the ester linkage is broken in the presence of water molecules and enzymes. Worksheet – 16 Polymerization
1
Polyamides, polyesters and polysaccharides are three types of condensation polymer.
a Complete the following tables, for three different polymerisation reactions A to C.
and
Reaction A
the monomers used for
the polymerization
structure of the polymer
represent carbon chains. formed
(show two units of each
monomer joined up)
the other product
that forms
name: ………………………………
type of polymer formed
(circle one)
polyamide
name of polymer
formed (circle one)
starch
synthetic or natural?
(circle one)
synthetic
polyester
nylon
soap
formula: ……………….
polysaccharide
terylene
sugar
natural
Reaction B
the monomers used
structure of the polymer
formed
(show two units of each
monomer joined up)
the other product that
forms
name: ………………………………
type of polymer formed
(circle one)
polyamide
name of polymer
formed (circle one)
starch
synthetic or natural?
(circle one)
synthetic
polyester
nylon
soap
formula: ……………….
polysaccharide
terylene
sugar
natural
Reaction C
the monomers used
structure of the polymer
formed
(show two units of each
monomer joined up)
the other product that
forms
name: ……………………………
formula: ………………..
type of polymer formed
(circle one)
polyamide
name of polymer
formed (circle one)
starch
synthetic or natural?
(circle one)
synthetic
polyester
nylon
soap
polysaccharide
terylene
sugar
natural
b What happens during condensation polymerisation? .................................................................................................................................................... ...................................................................................................................................................... c i Which of the three polymers formed in reactions A – C contains a linkage like that found in proteins? ............................................................................................................................................... ii What is the main difference between the structure of this polymer, and the structure of proteins? .............................................................................................................................................. ............................................................................................................................................... 2 a Only one of these molecules can be used to make a condensation polymer. Which one? ... i ii b Explain why the other molecule is unable to form a condensation polymer. ..................................................................................................................................... ...................................................................................................................................... ........................................................................................................................................ .