THE PERIODIC TABLE

The periodic table is a table of elements arranged in order of increasing proton number to show the similarities of the chemical elements with related electronic structures.

GROUPS

A vertical column of the periodic table containing elements with similar properties with the same number of electrons in the outer shell. They have an increasing number of inner shell as you descend the group.

There are 8 groups of elements.

The first column is called group 1 (alkali metals); the second group 2 (alkaline earth metals); and so on up to group 7 (halogens). The final column in the periodic table is called group 0 (inert gases or noble gases)

PERIODS

Horizontal rows of the periodic table. Within a period the atoms of all the elements have the same number of occupied shells but have an increasing number of electrons in the outer shell.

Between groups 2 and 3 is the block of elements known as transition elements

The periodic table can be divided into two by a blod line that starts beneath boron.

The elements to the left of this line are metals and those on the right are non metals.

The elements which lie on the line are called metalloids. These elements behave in some ways as a metal and in others as non metals.

GROUP 1 - the alkali metals

Consists of five metals lithium, sodium, potassium, rubidium, caecium and the radioactive element francium

lithium, sodium and potassium have the following properties:

They;

- are very reactive metals

- are good conductors of heat and electricity

- are soft metals, lithium is the hardest and potassium the softest.

- have low densities

-  have shiny surfaces when freshly cut with a knife

- have low melting points

- burn in oxigen or air with characteristic flame colours

- react vigorously with water to give an alkali solution of the metal hydroxide plus hydrogen gas. Potassium is the most reactive followed by sodium and then lithium

- react vigorously with halogens such as chlorine to form metal halides eg. Sodium chloride 

GROUP 2 - the alkaline earth metals

Consists of the five metals beryllium, magnesium, calcium, strontium and barium and the radioactive element radium.

Magnesium and calcium have the following properties.

- they are harder than those in group 1

- they are sivery-grey in colour when pure and clean. They tarnish quickly, however, when left in air due to the formation of metal oxides on their surfaces 

- they are good conductors of heat and electricity

- they burn in oxygen or air with characteristic flame colours to form solid white oxides

- they react with water but do so much less vigorously than the elements in group 1

GROUP 7 - the halogens

Consists of the four elements flourine, chlorine, bromine and iodine, and the radioactive element astatine.

chlorine, bromine and iodine have the following properties:

- they are coloured and darken going down the group

- they exist as diatomic molecules eg. Cl2, Br2 and I2.

- they show a gradual change from a gas (Cl2), through a liquid (Br2), to a solid (I2).

- they form molecular compounds with other non-metallic elements eg. HCl

- they react with hydrogen halides, which dissolve in water to form acidic solutions     - they react with metals to produce ionic metal halides eg. Iron (III) chloride

Colours of some halodens.

Chlorine - Pale green

Bromine - Red-brown

Iodine - Purple-black

Displacement Reaction

If chlorine is bubbled into a solution of potassium iodide, the less reactive halogen is displaced by the more reactive chlorine

 2KI + Cl2 = 2KCl + I2

The halogens have many varied uses

·         Floride is used in the form of florides in  water and toothpastes it reduces tooth  decay by hardening the enamel

·         Chlorine is used to make PVC plastic as well as in household bleaches. It is also used to kill bacteria and viruses in drinking water

·         Bromine is used to make disinfectants, medicines and fire retardants

·         Iodine is used in medicines and disinfectants and also as a photographic chemical

GROUP 0 - the noble gases

Helium, neon, argon, krypton, xenon, and the radioactive element radon make up the noble gases

They are colourless gases

  

Paper 6: Alternative to Practical

This paper is designed to test candidates’ familiarity with laboratory practical procedure.

Questions may be set from the following experimental contexts:
• simple quantitative experiments involving the measurement of volumes;
• speeds of reaction;
• measurement of temperature based on a thermometer with 1 °C graduations;
• problems of an investigatory nature, possibly including suitable organic compounds;
• simple paper chromatography;
• filtration;
• identification of ions and gases as specified in the Core curriculum.

Questions may be set requiring candidates to do the following:
• record readings from diagrams of apparatus
• describe, explain or comment on experimental arrangements and techniques
• complete tables of data
• draw conclusions from information given
• interpret and evaluate observations and experimental data
• describe tests for gases and ions, and/or draw conclusions from such tests
• plot graphs and/or interpret graphical information
• identify sources of error and suggest possible improvements in procedures
• plan an investigation, including suggesting suitable techniques and apparatus.

 Source
Cambridge IGCSE Chemistry
Syllabus code 0620
For examination in June and November 2011

Metals.

Properties Of Metals

1. Usually solid,occasionally liquid
2. Malleable
3. Ductile
4. Solid metals are shiny
5. Usually high melty and boiling points
6. Usually high density
7. Good thermal and electrical conductors.

Though metals share certain properties they differ in other ways.

ACIDS, BASES & SALTS

Acids & Alkalis

The word acid means 'sour' and all acids possess this property.

They are also soluble in water and are corrosive.

Alkalis are the chemical 'opposite' of acids.

They will remove the sharp taste from the acid.

They have a soapy feel.

Indicators are used to identify acids and alkalis.

They change colour when added to acids or alkalis.

Many indicators are dyes which have been extracted from natural resources

For example Litmus is a purple dye extracted from Lichens.

Universal Indicator is used to determine how acidic or alkaline a substance is. It is a mixture of many indicators.

The colours shown by this indicator can be matched against a pH scale. The pH scale runs from 0-14.

• A substance with a pH less than 7 is an acid. E.g HCL, Lemon Juice, Milk etc

• A substance with a pH greater than 7 is an alkaline. E.g toothpaste, bleach, Sodium Hydroxide etc

• One with a pH of 7 is neither acid nor alkaline, it is neutral e.g water, salt solution

A pH meter can also be used to determine the pH of a substance.

Bronsted-Lowery Theory

• An acids is an H+ ion (or proton) donor

• A base is an H+ ion (or proton) acceptor

The theory explains why a pure acid behaves differently from its aqueous solution, since for an acid to behave as an H+ ion donor it must have another substance present to accept the H+ ion. So the water, in the aqueous acid solution is behaving as a Bronsted-Lowery base and accepting H+ ion.

Amphoteric 

Is a substance which can behave both as a Bronsted-Lowery acid and as a Bronsted-Lowery base e.g water.

The Relative Strengths of Acids & Bases
Strong  Acids

• The relative strength of an acid is found by companing one acid with another.

• The strength of any acid depends upon how many molecules dissociate (or ionise) when the acid is dissolved in water.

 

• For hydrochloric acid all the hydrogen chloride molecules break up to form  H+ ions and Cl- ions.

• Any acid that behaves this way is termed a strong acid.

• They are good conductors of electricity,React quickly with metals, bases and metal carbonates.

• Sulfuric acid and nitric acid are also strong acids.

• A strong acid is still a strong acid even when it is in dilute solution.

 Weak Acids

• Weak acids such as ethanoic acid found in vinegar produce few hydrogen ions when it dissolves in water compared with a strong acid  of the same concentration.

• It is only partially ionised.

• Citric acid found in oranges and lemons, carbonic acid found in soft drinks, sulfurous acid (acid rain) are example of weak acids.

• They are poor conductors of electricity and have slower reactions with metals,bases,and metal carbonates

The concentration of an acid tells you the proportion of water and acid present in aqueous solution.


To say an acid is strong does not mean it is concentrated.


A strong acid is strong even when it is in dilute solution.


A weak  acid is weak even when it is concentrated.



Strong  Bases

• The strength of a base is found by comparing one base to another and it is dependent upon the dissociation of the base in aqueous solution.

• An Alkali:- is a base which produces hydroxide ions,OH^-(aq), when dissolved in water.

• NaOH is a strong alkali because when it dissolves in water its lattice breaks up completely to produce ions.

    

•  another example is potassium hydroxide

Weak Bases

A weak alkali, such as ammonia, produces fewer hydroxide ions when it dissolves in water. it is only partially ionises.

Neutralising an acid

• If you wish to neutralise an acid such as HCl you can use an alkali such as NaOH.

• If equal volumes of the same concentration of HCl and NaOH are added to one another, the resulting solution is found to have a pH of 7.

• The acid has been neutralised and a neutral solution has been formed.

Formation of  salts

• A normal salt is formed when all the hydrogen ion of an acid have been replaced by metal ions or by ammonium ion. 
• Normal salts can be grouped into those which are soluble in water or those which are insoluble in water.
• The following salts are soluble in cold water
• all nitrates
• all common sodium, potassium and ammonium salts
• all chlorides except lead, silver and mercury
• all sulfates except lead, barium and calcium

METHODS OF PREPARING  SOLUBLE SALTS 

1. Acid + Metal 

• Can only  be used with the less reactive metals.
• It will be very dangerous to use a reactive metal such as Na.
• The metals usually used are the MAZIT. ie. magnesium, aluminium, zinc, iron and tin.

         Typical Experiment

• Excess Mg ribbon is added to dilute nitric acid.
• During the addition an effervescence is observed due to the production of hydrogen gas. The hydrogen ions from the nitric acid gain electrons from the metal atoms to produce hydrogen gas.

• The excess magnesium is removed by filtration.
• The Mg(NO3 )2 solution is evaporated slowly to form a saturated solution of the salt.
• The hot concentrated Mg(NO3 )2 solution produced is tested by dipping a cold glass rod into it, if the salt crystals form at the end of the rod the solution is ready to crystallise and left to cool
• Any crystals produced on cooling are filtered and dried between tissues. 

 2. Acid + Carbonate

This method can be used with any metal carbonate and any acid.

     Experiment

•      Copper (II) carbonate added in excess to dilute nitric acid
•      Effervescence observed due to production of CO2 

     Equation  
 

• Metal carbonates contain carbonate ions in this reaction, the carbonate ions react with H+ ions in the acid.

 

  3. ACID + ALKALI (Soluble Base)

• This method is used to prepare salts of very reactive metals such as potassium and Na
• It will be too dangerous to add the metal directly to the acid.
• We solve the problem by using an alkali which contains the particular reactive metal whose salt we wish to prepare.

• A base is a substance which neutralises an acid producing salt and water.
• If the base is soluble,it is termed alkali
• Most metal oxides and hydroxides (as well as ammonia solution) are bases.

 
    4. ACID + INSOLUBLE BASE

• This method can be used to prepare a salt of an nonreactive metal such as Pb or Cu.
• The acid is neutralised using the particular metal oxide.
•  Some warming up of the reactants may be necessary
• Example is the neutralisation of  sulfuric acid by copper (II) oxide to produce  copper (II) sulfate
• After slow evaporation to concentrate the solution the solution is left to crystallise. Crystals of copper (II) sulfate are produced.    

METHODS OF PREPARING INSOLUBLE SALTS

An insoluble salt, such as barium sulphate can be made precipitation. In this case, solutions of the two chosen soluble salt are mixed to produce barium sulphate, barium chloride and sodium sulphate can be used. The barium sulphate precipitate can be filtered off, washed with distilled water and dried.

It should be noted that salts like barium sulphate dissolve to a very small extent, for example, 1 litre of water will dissolve 2.2*10-3g of barium sulphate at 25oC. This substance and substances like it said to be sparingly soluble.

 

It should be noted that salts like barium sulphate dissolve to a very small extent, for example, 1 litre of water will dissolve 2.2*10-3g of barium sulphate at 25oC. This substance and substances like it said to be sparingly soluble.

Titration

• Is a method of volumetric analysis in which a volume of an acid is added to a known volume of another an alkali slowly from a burette until an  end part is reached. 
• If an acid and alkali are used, then an indicator is used to show that the end-part has been reached.

In the laboratory, the titration of HCl with NaOH is carried out in the following way

1. 1.       25cm³     of NaOH solution is pipette into a conical flask to which a few drops of phenolphthalein indicator have been added. Phenolphthalein is pink in alkaline conditions but colorless in acid.

1. 2.      A 0.10 moldm^-3 solution of hydrochloric acid is placed in the burette using a filter funnel until it is filled up exactly to the zero mark

1. 3.      The filter funnel is now removed

1. 4.      The HCL is added to the NaOH solution in small quantities the contents of the flask must be swirled after each addition of acid fir through mixing

1. 5.      The acid is added until the alkali has been neutralized completely. This is shown by the pink colour of  the indicator just disappearing.

1. 6.      The final reading on the burette at the end-point is recorded and further titration carried out until consistent results are obtained.

 Questions

1. 1.24.2cm³ of a solution containing 0.20moldm^-3 of HCL just neutralized 25.0cm³ a potassium hydroxide solution. What is the conclusion of this KOH? 
2. 22.4cm³ of a solution containing 0.10mol dm^-3 of H₂SO₄ just neutralized 25.0cm^-3 of a NaOH solution. What is the conclusion .what is the conclusion of this NaOH solution.