Presentation College Chemistry- Mr Adrian Ramlal's Class!: 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 Na₂Cl_5.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 SO₄^2- 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 (SO₄^2-), carbonate (CO₃^2-), nitrate (NO₃^-) and ammonium (NH₄⁺).

The Charges on some common negative ions (anions)

1- ions	2- ions	3- ions
Fluoride, F^-	Sulphate, SO₄^2-	Phosphate, PO₄^3-
Chloride, Cl^-	Sulphite, SO₃^2-	Nitride, N^3-
Iodide, I^-	Carbonate, CO₃^2-
Hydroxide, OH^-	Oxide, O^2-
Manganate (VII), MnO₄^-	Sulphide, S^2-
Hydrogencarbonate, HCO₃^-	Chromate (VI), CrO₄^2-
Ethanoate, CH₃COO^-	Dichromate (VI), Cr₂O₇^2-
Methanoate, HCOO^-	Ethanedioate, C2O4^2-
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-

Composition of air (% by volume):

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)

Wednesday, May 30, 2012

Ionic Equations