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)
No comments:
Post a Comment