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Presentation College Chemistry- Mr Adrian Ramlal's Class!
This is your key to EXAM SUCCESS!!!
Wednesday, May 30, 2012
FORM 3 TERM
III Class Notes and Revision Manual
TERM 3 ( 12 weeks) Mr Ramlals Classes
BALANCED
CHEMICAL COMPOUNDS AND EQUATIONS:
- Recall formation of ionic and covalent
compounds (stable octet).
- Must be able to write balanced compounds.
- What are chemical equations?
- Must be able to write chemical equations.
- Must be able to balance chemical
equations.
- Derive ionic equations.
- Show knowledge of the types of chemical
equations
THE
PERIODIC TABLE:
·
Who
is credited for the periodic table?
·
How
are elements placed in the periodic table?
·
Discuss
periods and groups.
·
Discuss
trends in atomic radius and thus reactivity in relation to periods and groups.
·
Define
electronegativity and its trend across the periodic table.
REDOX
CHEMISTRY : OXIDATION AND REDUCTION
- What are oxidation state numbers? Note
anomalies
- Discuss changes in oxidation state numbers
in relation to gaining and losing of electrons.
- Must be able to define oxidation and
reduction in relation to loss and gain of electrons.
TOPIC TO BE CONTINUED IN FORM 4
WITH RELEVANT SBA LABS
TOPIC AS
OUTLINED IN SEMP SYLLABUS:
- Air-composition
- Air-pollution; (project on sources of air pollution, effects of
air pollution globally and locally, factors which affect the implementation
of policies for reducing pollution to the atmosphere.
Notes:
BALANCE CHEMICAL COMPOUNDS AND EQUATIONS:
- Recall formation of ionic and covalent
compounds (stable octet).
The
majority of atoms do not exist as single entities. Instead they bond together
to form groups called molecules. The
only atoms that exist singly are the members of group 8 in the Periodic Table.
These are called the noble or inert gases. The noble gases do not take part in
ionic or covalent bonding. The need for bonding may be explained by looking
at the electronic configuration of these gases.
He-2
Ne-2,8
Ar-2,8,8
Kr-2,8,18,8
The
electronic configuration of each noble gas shows it has 8 electrons in its
outermost electron shell. For atoms this is a stable configuration. All other atoms
seek to achieve a similar configuration. They do so by bonding.
There are
three types of inter-molecular bonding-
1) Ionic 2)Covalent 3)
Metallic
When an
atom bonds, only the electrons in its outer shell are affected. These are
called the valence electrons.
1) Ionic bonding- Ionic bonding takes
place between a metal and non-metal atoms. It involves the transfer of
electrons e.g. Sodium and Chlorine. Sodium loses an electron while chlorine
gains an electron.
Ions: Ions
are charged atoms and molecules. The charge on an ion is also called its
valency. There are two types of ions:
Cations:
These are positive ions that result when a metal atom loses electrons.
Anions:
These are negative ions that result when a non-metal atom gains electrons. The
numerical value of the valency is equal to the number of electrons lost or
gained.
2) Covalent Bonding-
Covalent
Bonding involves the sharing of electrons. It occurs between non-metal atoms.
Atoms may be of the same time or of different types. In covalent bonding the
outer shell of each atom overlaps. Both atoms share the electron pair. In
covalent bonding, atoms are held together by their shared electron pairs.
Therefore
covalent compounds can be formed between two identical non-metal atoms.
- Must be able
to write balanced compounds.
Before even
beginning to write a chemical equation one must know how to balance a chemical
compound. It is the first step to writing balanced equations. What this
essentially means is that we must know how to write in symbols the correct
ratio of elements in a compound, for instance why do we write NaCl and not Na2Cl
or even Na2Cl5. The simplest way to create these
compounds is by knowing their charges and combining them to create them
electrically neutral.
Molecular
formulae-
The
molecular formulae of a compound gives:
- the number of atoms of each element
present in one molecule for covalent compunds
- the ratio of ions present in one formula
unit for ionic compounds (because the term “molecule” cannot be applied to
ionic compounds)
The
molecular formula of a compound gives the actual numbers of the different types
of atoms in one molecule of a covalent compound or the ratio of ions present in
one formula unit of an ionic compound.
Eg. A) The
molecular formula of glucose is C6H12O6. This
means that 1 molecule of glucose contains:
- 6 atoms of carbon
- 12 atoms of hydrogen
- 6 atoms of oxygen
b) the
molecular formula of ammonium suphate is (NH4)2SO4
The ratio
of ions is:
- Two NH4+ ions
- One SO42- ions.
Note:
chemical formulae are really worked out by experiment however here you will
learn to use charges and simple tools to help you write the molecular formulae
of a compound correctly using other types of information. You need to write formulae correctly in order to write equations.
Working out numbers and Ratios in compounds-
We can work
out the formulae of compounds by using charges or oxidation numbers.
Using
Charges-
This method
only works for ionic compounds. The two tables below contain lists of the
charges carried by commom anions and cations (including polyatomic ions). A
polyatomic ion (sometimes called a radical) is a charged group of atoms that often
occur together in compounds. Examples are sulphate (SO42-),
carbonate (CO32-), nitrate (NO3-)
and ammonium (NH4+).
The Charges
on some common negative ions (anions)
1- ions
|
2- ions
|
3- ions
|
Fluoride, F-
|
Sulphate, SO42-
|
Phosphate, PO43-
|
Chloride, Cl-
|
Sulphite, SO32-
|
Nitride, N3-
|
Iodide, I-
|
Carbonate, CO32-
|
|
Hydroxide, OH-
|
Oxide, O2-
|
|
Manganate (VII), MnO4-
|
Sulphide, S2-
|
|
Hydrogencarbonate, HCO3-
|
Chromate (VI), CrO42-
|
|
Ethanoate, CH3COO-
|
Dichromate (VI), Cr2O72-
|
|
Methanoate, HCOO-
|
Ethanedioate, C2O42-
|
|
Bromide, Br-
|
|
|
The Charges
on Some common positive ions (cations)
1+ ions
|
2+ ions
|
3+ ions
|
Metals of Group 1: Li+, Na+, K+
|
Metals of Groups II: Mg2+, Ca2+,
Ba2+
|
Aluminum, Al3+
|
Hydrogen, H+
|
Lead(II), Pb2+
|
Iron(III), Fe3+
|
Ammonium, NH4+
|
Iron (II), Fe2+
|
Chromium (III), Cr3+
|
Silver, Ag+
|
Zinc, Zn2+
|
|
Copper (I), Cu+
|
Tin (II), Sn2+
|
|
|
Copper(II), Cu2+
|
|
|
|
|
Here is how
you can work out formulae using the charges on the ions in an ionic compound.
Step 1:
Write the name of the compound
Step 2: Write
down the symbol of the elements (or polyatomic groups) and their charges.
Step 3:
Since the total negative charge must equal the total positive charge, balance
out the charges by adjusting the numbers of ions as necessary. Do not change
the charge on any of the ions.
Step 4:
write the formula using the numbers of each ion as subscript.
Here are
some examples:
Step 1:
Calcium Oxide
Step 2: Ca
2+ O 2-
Step 3:
These are balanced
Step 4: Ca1
O1
Note: This
formula is written as CaO since the ratio of ions is 1:1. Where the number of
an atom is 1, the 1 is never written down-it is understood.
Step 1:
Zinc Chloride
Step 2:
Zn2+ Cl-
Step 3:
Zn2+ gives 2+ Cl- ; Cl- gives 2-
Step 4:
ZnCl2
Step 1:
Aluminum Sulphate
Step 2: Al
3+ SO4 2-
Step 3:
Al3+, Al3+ gives 6+ SO42- ; SO42-, SO42- gives 6-
Step 4:
Al2(SO4)3
The formula
is Al2(SO4)3
Note: the
subscript 3 is written outside the brackets to indicate three sulphate groups.
USING
OXIDATION NUMBERS:
This method
can be used for all compounds.
Each
element in a compound can be given an oxidation number. An oxidation number is
assigned to an element by a set of rules.
If you are
given the oxidation numbers of the elements in any compound, you can work out
the formula in much the same way as when you use charges. The sum of the
oxidation numbers of elements in a compound is zero.
For some
elements, the oxidation number is always or almost always the same. However,
some elements, such as iron, have variable oxidation numbers. Roman numerals
are used to distinguish between the compounds of such elements. Iron (II)
Chloride and Iron (III) chloride are the names of Iron compounds of oxidation
numbers +2 and +3 Respectively.
Here is the
method you need to use:
Step 1:
Write the name of the compound
Step 2:
Write the elements in the compound and their oxidation numbers
Step 3:
Balance out the numbers
Step 4:
Write the formula.
Here are
some examples:
Step 1:
Hydrogen Chloride
Step 2 H+1
Cl-1
Step 3: +1
and -1 add up to zero
Step 4: HCl
The formula
is HCl
Step 1:
Sulphur (VI) Oxide
Step 2: S+6
O-2
Step 3: S+6
O-2, O-2;+6 and -6 add up to zero.
Step 4: SO3
The formula
is SO3
Step 1:
Sulphur (IV) Dioxide
Step 2: S+4
O-2
Step 3: S+4
O-2; +4 and -4 add up to zero
Step 4: SO2
The formula
is SO2
- What are
chemical equations?
- Must be able
to write chemical equations.
- Must be able
to balance chemical equations.
Chemical
Equations and Reactions-
Substances
can be changed in several ways eg, heating and mixing with other substances.
Two types of changes can occur:
1) Physical Changes
2) Chemical Changes
Physical
changes are those in which the chemical formulae of the substances remain the
same. Some physical changes are dissolving, melting, freezing and forming a
mixture. When ice melts, this is a physical change because water and ice have
the same chemical formula, H2O. Physical changes are also easily reversible.
Water can be easily reconverted to ice.
Chemical
changes are those in which the products or end substances are totally different
from the reactants or the starting materals. Changes in chemical formulae
occur. A chemical change is also referred to as a chemical reaction.
When
Magnesium and oxygen are heated together, a chemical reaction occurs, because a
totally new substance, Magnesium Oxide is formed. Chemical changes are not
usually easily reversible.
Signs that
a chemical reaction is occurring may include:
- A colour change
- A change in temperature. The mixture
becomes hot or cold.
- Effervescence (a gas is given off)
Chemical
Equations-
Just as the
formula is the shorthand way of writing the name of a compound, chemical
equation is the shorthand way of representing a chemical reaction. A chemical
equation consists of:
Reactants
---------à
Products
The
following steps can be used to write a balanced equation.
The
following steps can be used to write a balanced chemical equation.
1) Write the word equation for the
reaction. This consists of the names of the reactants and the products.
Magnesium +
Hydrochloric Acid----à Magnesium Chloride and Hydrogen
2) Convert the words into symbols and
formulae
Mg +
HCl------à
MgCl2 + H2
3) Balance the equation so that there
are the same numbers of atoms and ions of each element on both sides of the
equation. Begin by underlining each formula. DO NOT change the numbers inside
the boxes, since this will change the chemical formulae. Look at the numbers of
elements on each side and balance them one at a time by putting simple whole
numbers at the left of each underlined compound.
Mg + HCl ------à MgCl2 + H2
There is 1
reactant H and 2 Product H. To balance them put a 2 outside HCl.
Mg + 2HCl ----------à MgCl2 + H2
All the
numbers balance so the equation is:
Mg + 2HCl
----------à MgCl2 + H2
4) Write the state symbols for each
substance in the equation. The state symbol tells the physical state of the
substance during the reaction. There are four state symbols
s- solid aq-aqueous
l-liquid g-gas
state
symbols are put in brackets at the bottom right of the substances.
Eg. Mg(s) + 2HCl(aq) ------à MgCl2(aq) + H2(g)
This is
your complete equation.
Other
examples:
1) Word Equation: Magnesium + Oxygen-à Magnesium Oxide
Convert to
symbols: Mg + O2 -----à
MgO
Balance
Oxygens: Mg + O2 ----à
2MgO
Balance
Magnesium: 2Mg + O2
----à 2MgO
Put in
state symbols: 2Mg(s) + O2
--à 2MgO (s)
2) Word Equation : Aluminum + Sulphuric
Acid --à
Aluminum Sulphate + Hydrogen
Convert to
symbols:
Convert to
symbols: Al + H2SO4 ---à Al2(SO4)3 + H2
Balance
Aluminum: 2 Al + H2SO4 ----à Al2(SO4)3 + H2
Balance
Sulphate: 2 Al + 3 H2SO4 ----à Al2(SO4)3
+ H2
Balance
Hydrogen: 2Al + 3H2SO4 --à Al2(SO4)3 + 3H2
Put in
state symbols: 2Al(s) = 3H2SO4(aq) --à Al2(SO4)3(aq) + 3H2(g)
Some tips
when writing equations:
1) When writing formulae, remember
gaseous elements exist in their natural state as diatomic molecules.
Hydrogen-
H2, Fluorine-F2, Chlorine- Cl2, Nitrogen-N2, Iodine-I2, Bromine-Br2, Oxygen-O2
2) When balancing the equation, start
by balancing elements which occur only in one formula on both sides of the
equation, eg.
Ca + H2O --à Ca(OH)2 + H2
Start with
O since this occurs only in H2O and Ca(OH)2
Tips for
writing Ionic Equations-
1) Elements always have a valency of 0
2) Covalent compounds cannot be broken
into ions , eg. H2O, CO2 and NH3.
3) Two Ions that come from the same
compound on the same side of the equation are rejoined in the final ionic
equation.
Writing
Ionic Equations-
Example #1: Write the balanced net ionic equation for
the reaction of aqueous sodium hydroxide and aqueous hydrochloric acid
Step #1: Write the balanced GENERAL
EQUATION - In order to write this equation, you must decide what the
products are. This example problem is an acid-base reaction. The products will be a salt (NaCl) and
water. After you have written the reaction, it must be balanced.
Step #2: Write the TOTAL IONIC EQUATION -
Here, each reactant and product is studied to determine whether it dissociates
in solution. If it is a strong electrolyte, it is written as ions. If it isn't
a strong electrolyte it is written as a molecule.
Because NaOH, HCl and NaCl are strong electrolytes they are
written as ions. Water is a nonelectrolyte and should be written as a molecule.
Step #3: Write the NET IONIC EQUATION - Each
species that does not undergo a change is called a "spectator ion".
These species are removed from the equation leaving the balanced net ionic
equation
In this example, Na+ and Cl- are spectator ions. They do not
undergo change in the reaction. Therefore, they are removed.
Types of Chemical Reactions-
There are many types of chemical reactions.
Some of these are-
1) Combination Reactions- This is when
two reactants combine to give one product.
S(s) +
O2(g) ----à
SO2(g)
Na2O(g) + CO2(g) ------à
Na2CO3(s)
2) Decomposition- When some substances are heated, they can
break down or decompose.
2NaNO3(s) -----à 2NaNO2(s) + O2(g)
3) Displacement- A metal or a non-metal
may displace another from its salt.
Zn(s) +
CuSO4(aq) ---à
ZnSO4 (aq) + Cu(s)
Cl2(g) + 2KI(aq)
----à
2KCl(aq) + I2(aq)
4) Double Decomposition- When salt
solutions are mixed together, they exchange ions.
2NaNO3(aq) + ZnSO4(aq)
---à Zn(NO3)2(aq)
+ Na2SO4(aq)
5) Neutralisation- When acids and bases
react until their pH is exactly 7 or neutral.
HCl (aq) +
NaOH (aq) -----à
NaCl(aq) + H2O(l)
Exam
Structure and Topics-
Section A Consists
of 30 Simple Multiple Choice Questions from Terms 1, 2 and 3
Section B
Consists of 5 Structured questions from these topics
1) Balancing Equations (You will need
to know how to form compounds as you will only be given the word equations) -15 marks
2) Writing Balanced Ionic Equations-15
marks
3) Atomic Structure, Charge Tables and
Ions/Covalent bonding (dot cross diagram knowledge necessary)/Properties of
metals and their structure
4) Separation Techniques-
Distillation/Solutions
5) Air, Air Composition, Greenhouse
gases, Consequences of Greenhouse gases
With Regard
to question 5- The air we breathe is composed of many gases, the main ones are-
Nitrogen (N2): 78.09%
Oxygen (O2): 20.95%
Argon (Ar): 0.93%
Carbon dioxide (CO2): 0.038%
Others (less than 0.002% each): Neon (Ne), Helium (He),
Krypton (Kr), Hydrogen (H2), Xenon (Xe).
A greenhouse
gas (sometimes abbreviated GHG) is a gas in an atmosphere that absorbs and emits
radiation within the thermal infrared range. This process is the fundamental
cause of the greenhouse effect.[1] The primary greenhouse gases in the Earth's
atmosphere are water vapour, carbon dioxide, methane, nitrous oxide, and ozone. In the Solar System, the atmospheres of Venus, Mars, and Titan also contain gases that cause greenhouse
effects. Greenhouse gases greatly affect the temperature of the Earth; without them, Earth's surface would be on
average about 33 °C (59 °F)[note 1]
colder than at present.[2][3][4]
However, since
the beginning of the Industrial
Revolution, the
burning of fossil fuels has contributed to the increase in carbon
dioxide in the atmosphere from 280 ppm to 390 ppm, despite the uptake of a
large portion of the emissions through various natural "sinks"
involved in the carbon cycle.[5][6] Anthropogenic carbon dioxide (CO2 )
emissions (i.e., emissions produced by human activities) come from combustion of carbonaceous fuels, principally wood, coal, oil, and natural gas.[7]
Greenhouse
gases
Atmospheric
absorption and scattering at different electromagnetic wavelengths. The largest absorption band of carbon dioxide is in the infrared.
Greenhouse gases
are those that can absorb and emit infrared
radiation.[1] In order, the most abundant greenhouse gases in
Earth's atmosphere are:
- water vapor (H2O)carbon dioxide (CO2)methane (CH4)nitrous oxide (N2O)ozone (O3)
Atmospheric
concentrations of greenhouse gases are determined by the balance between sources
(emissions of the gas from human activities and natural systems) and sinks (the
removal of the gas from the atmosphere by conversion to a different chemical
compound).[8] The proportion of an emission remaining in the
atmosphere after a specified time is the "Airborne fraction" (AF). More precisely, the annual AF
is the ratio of the atmospheric increase in a given year to that year’s total
emissions. For CO2 the AF over the last 50 years (1956–2006) has
been increasing at 0.25 ± 0.21%/year.[
Further notes and hints will be uploaded to
Preschem.blogspot.com (on wed 30 may 2012)
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