topic 4 periodic table

� INTRODUCTION

Hello and letÊs start Topic 4! In this topic, you will learn about the periodic table. Before we go further, do you know that there are 118 discovered elements in nature? Most of these elements are naturally occurring elements. However, a few of these elements are made up artificially in nuclear reactors. Elements with the same chemical properties were grouped together by chemists, resulting in the development of the periodic table. This systematic method of classifying elements has enabled us to study and generalise the chemical and physical properties of elements in the same group. �We will learn more about the periodic table as we track back its history and study how the groups and periods of the periodic table can be analysed. This is followed by the electronic structures and the periodic table, and properties and usages of transition elements. �

LEARNING OUTCOMES

By the end of this topic, you should be able to:

1. Analyse the periodic table;

2. Summarise the electronic structures and periodic table;

3. Identify properties and usages of transition elements;

4. Identify the electronic structure, group trends, physical propertiesand chemical properties of Group 1 and Group 17;

5. Summarise noble gases; and

6. Identify the properties and classification of Period 3 elements.

TTooppiicc

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� Periodic Table

TOPIC 4 PERIODIC TABLE �

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Then, we will examine the electronic structure, group trends, physical properties and chemical properties of Group 1 and Group 17. Last but not least, we will look at noble gases and Period 3 elements. Are you ready now? Let us start the journey!

HISTORY OF THE PERIODIC TABLE

Let us now review the history of the periodic table as well as the events which led to the development of the modern periodic table. Do you know that the majority of the elements that we know today were actually discovered during the 18th and 19th century? You will notice that elements with similar properties were grouped together systematically in a table. This marked the beginning of the development of the periodic table. Chemists such as Lavoiser, Dobereiner, Newlands, Meyer, Mendelev and Mosely contributed to the development of the periodic table in use today. We will now read about their respective contributions.

4.1.1 Antoine Lavoisier (1743–1794)

Do you know that Antoine Lavoisier (Figure 4.1) was the first scientist to classify elements into four groups? He classified substances, including light and heat, into metals and non-metals.

Figure 4.1: Antoine Lavoisier (1743�1794) Source: http://www.sciencephoto.com

However, his classification was not successful due to wrong information. For example, non-elements such as heat and light, and compounds such as silica, magnesia, chalk, barita and alumina were included in his classification table.

4.1

� TOPIC 4 PERIODIC TABLE 76

Table 4.1: Antoine LavoisierÊs 1789 Classification of Substances

Acid-making Gas-like Elements Metallic Elements Earthy Elements

Sulphur Light Cobalt, Mercury, Tin

Lime (Calcium Oxide)

Phosphorus Caloric (Heat) Copper, Nickel, Iron Magnesia (Magnesium Oxide)

Charcoal (Carbon) Oxygen Gold, Lead, Silver, Zinc

Barytes (Barium Sulphate)

Azote (Nitrogen) Manganese, Tungsten

Argilla (Aluminium Oxide)

Hydrogen Platina (Platinum) Silvex (Silicon Dioxide)

Source: http://www.docbrown.info/page12/gifs/Lavoisier1789.gif

4.1.2 Johann Dobereiner (1780–1849)

Johann Dobereiner (Figure 4.2) divided the elements into groups. Each group consists of three elements with similar chemical properties and is called a triad. In each triad, the atomic weight of the middle element is the average of the other two elements. According to the Law of Triad, the atomic mass of sodium is the mean of the total atomic mass of lithium and potassium. Thus, the atomic mass of sodium is 23 (refer to Table 4.2).

Figure 4.2: Johann Dobereiner (1780�1849) Source: http://elements-table.com/history/


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Table 4.2: Law of Triad

Element Symbol A (Atomic Mass)

Lithium Li 7

Sodium Na 23

Potassium K 39

Mean of Li + K = (7 + 39)/2 = 46/2 = 23 (The value of Na) However, this classification was unsuccessful because the classification was limited to a few elements only. Then, other scientists realised that there was a relationship between the properties and atomic masses of the elements, as shown in Table 4.3.

Table 4.3: Relationship between the Properties and Atomic Masses of the Elements

Triads IIII IIV

Elements Copper Cu

Silver Ag

Gold Au

Zinc Zn

Cadmium Cd

Mercury Hg

Atomic weights 63�5 108 197 65 112.5 200

Mean Weights � 130.25 � � 132.5 �

Source: http://www.tutornext.com/ws/402-g-limit

4.1.3 John Newlands (1837–1898)

Another chemist that contributed to the existence of the periodic table was John Newlands (Figure 4.3).

Figure 4.3: John Newlands (1837�1898) Source: http://elements-table.com


Newlands arranged all the known elements horizontally in the ascending order of their atomic masses. Each row consisted of seven elements. He found that elements with similar properties recurred at every eighth element. This arrangement was known as the LLaw of Octaves. �However, this law was only obeyed by the first 17 elements. Thus, it was not successful. There were no positions allocated for elements yet to be discovered. However, Newlands contibution to the development of the periodic table was very important as he was the first chemist who discovered the existence of periodicity in the elements.

4.1.4 Lothar Meyer (1830–1895)

Lothar Meyer (Figure 4.4) plotted a graph of atomic volume against atomic mass for all known elements. He found that elements with the same chemical properties occupied the same relative positions on the curve. He showed that the properties of the elements were in a periodic pattern with their atomic masses. Hence, Meyer also proved that the properties of the elements recur periodically.

Figure 4.4: John Newlands (1837�1898) Source: http://www.wou.edu


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4.1.5 Dmitri Mendeleev (1834–1907)

Dmitri Mendeleev (Figure 4.5) showed that the properties of elements changed periodically with their atomic mass. He arranged the elements in the order of increasing atomic mass and grouped them according to similar chemical properties. He was able to predict the properties of undiscovered elements and left gap for these elements.

Figure 4.5: Dmitri Mendeleev (1834�1907) Source: http://chemistry.about.com

Mendeleev had also correctly predicted the properties of the elements gallium, scandium and germanium which were only discovered much later. MendeleevÊs table was used as a blueprint for the modern periodic table. Figure 4.6 shows MendeleevÊs periodic table.

Figure 4.6: MendeleevÊs periodic table Source: http://www.msnucleus.org


4.1.6 Henry J. G. Moseley (1887–1915)

Henry J. G. Moseley (Figure 4.7) studied the x-ray spectrum of elements. He concluded that the proton numbers should be used as a basis for the periodic change of chemical properties instead of the atomic mass. He rearranged the elements in the ascending order of their proton numbers.

Figure 4.7: Henry J. G. Moseley (1887�1915) Source: http://en.wikipedia.org

Similar to Mendeleev, Mosely left gaps for elements yet to be discovered. He produced a periodic table which was almost the same as MendeleevÊs periodic table. Thus, he confirmed the work of Mendeleev. Due to MoseleyÊs work, the periodic table was successfully developed and being used today. The modern periodic table is based on the arrangement of elements in the ascending order of their proton numbers. Finally, the periodic table is as what we see today.


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4.1.7 Modern Periodic Table

Based on our earlier discussions about the early history of the periodic table, what can you conclude about it? How would you define the periodic table?

� Later, Glenn Seaborg (Figure 4.8) discovered that the transuranium elements have atomic numbers from 94 to 102, resulting in the redesign of the periodic table.

Figure 4.8: Gleen Seaborg Source: http://www.wired.com

Technically, both the lanthanide and actinide series of elements are to be placed between the alkaline earth metal and the transition metal. However, by doing this, the periodic table would be too wide. Thus, the lanthanide and actinide series of elements were placed under the rest of the periodic table. This is the periodic table that we use today. Dr Seaborg and his colleagues were also responsible for identifying more than 100 isotopes of elements.

The pperiodic table is a cclassification of elements whereby elements with the ssame chemical properties are placed in the ssame group. This makes the study of the chemistry of these elements eeasier and mmore systematic.


Figure 4.9 shows the modern periodic table. From here on, we will do an in-depth study of the periodic table. Based on calculation, there are 118 elements in the current periodic table but for the purpose of study for this module, only 111 elements will be considered.

Figure 4.9: The modern periodic table

Source: http://www.webelements.com/

1. List the name of the chemists who played a significant role in theearly development of the periodic table.

2. What was the conclusion of the study by Henry J. G. Moseley? 3. Define the periodic table in your own words. 4. Differentiate between the old version and the modern version of

the periodic table.

SELF-CHECK 4.1


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ANALYSIS OF THE PERIODIC TABLE

Let us refer to the periodic table to see the arrangement of the elements. The elements in the periodic table are arranged in rrows called the pperiods and columns which are known as the ggroups. Notice that in the periodic table, the atomic number increases when moving across a row or a period. Let us learn more about groups and periods.

4.2.1 Groups

Firstly, there are 18 groups of elements in the periodic table. Some of these groups have special names:

(a) Group 1 elements are called aalkali metals.

(b) Group 2 elements are called aalkaline earth metals.

(c) Group 3 to Group 12 elements are known as ttransition elements.

(d) Group 17 elements are called hhalogens.

(e) Group 18 elements are called nnoble gases. Do you know that there is a guideline for you to easily differentiate between elements in the group? The guideline is based on their classification as metals and non-metals:

4.2

1. Visit http://www.webchem.net/JSPT.htm and see the animatedversion of the periodic table. Based on the periodic table, findthe:

(a) Atomic weight, melting point, boiling point and electronconfiguration of sodium, magnesium, calcium andbromine; and

(b) Properties of other elements from the periodic table. 2. Compare and contrast the findings of each of the scientistsÊ

periodic table against the modern periodic table. What are thesimilarities and differences among the periodic tables?

ACTIVITY 4.1


(a) MMetals The elements in Groups 1 to 13 are metals.

(b) NNon-metals

The elements in Groups 15, 16 and 17 are non-metals. Carbon and silicon from Group 14 are also non-metals.

Keep in mind that each member of a group shows similar chemical properties although their physical properties such as density, melting point and colour show a gradual change when descending the group.

4.2.2 Periods

How about periods? What can you say about it? The hhorizontal rows in the periodic table are called pperiods. Let us take a look at the number of rows in the periodic table. You can see that there are seven rows from Period 1 to Period 7. The elements are arranged horizontally in the ascending order of their proton numbers in the periodic table. This means that when going across the rows from left to right, the proton number increases. How do we determine the position of the period of an element in the periodic table? The position of the period of an element in the periodic table is determined by the nnumber of shells occupied with electrons in the atom of the particular element. Do you know that the first three periods are called the short periods? Period 1 has two elements only � Hidrogen and Helium, while Periods 2 and 3 have eight elements each. As for Periods 4 and 5, they have eighteen elements each and they are called the long periods. Period 6 has thirty-two elements and not all the elements can be listed on the same horizontal row. The elements with proton number 58 to 71 are separated and are grouped below the perodic table. These elements are also known as the LLanthanide Series. Lastly, let us look at Period 7. Period 7 has thirty-one elements and not all the elements can be listed on the same horizontal row. The elements with proton number 90 to 103 are grouped below the perodic table. These elements are also known as the AActinide Series.


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1. Describe what you understand by the following terms. Providespecific examples for each term: (a) Groups of elements; and (b) Periods of elements.

2. Find the number of elements from the following periods by

referring to the periodic table. (a) Period 1; (b) Periods 2 and 3; (c) Periods 4 and 5; (d) Period 6; and (e) Period 7.

3. Fill the following particulars in the diagram below:

(a) Number the groups; (b) Number the periods; (c) Identify the position of alkali metals, alkaline earth metals,

transition elements, halogens and noble gases; and (d) Identify the location of lanthanides and actinides.

Source: http://www.chemeddl.org/collections/tsts/PeriodicTable.gif

SELF-CHECK 4.2


ELECTRONIC STRUCTURES AND THE PERIODIC TABLE

Before we go futher, let us recall what we have learnt in the previous topic on the subject of the electronic structure of an atom. The numbering of the shells starts from 1, 2, 3 and so on, starting from the one closest to the nucleus. The first shell can hold a maximum of two electrons, the second shell eight electrons and the third shell eight electrons if the number of electrons is less than 20. If the number of electrons is more than 20, the third shell can hold a maximum of up to 18 electrons. Do you still remember valence electrons? What are they? VValence electrons are electrons found in the ooutermost occupied shell of an atom. The number of valence electrons in an atom can be determined from its eelectronic structures. All members of the same group have the same number of valence electrons. The number of valence electrons of Group 1 and Group 2 elements is the same as its group number. For example, members of Group 1 have one valence electron each. For elements more than two valence electrons, namely, those in Group 13 to Group 18, this is how we determine them:

The group number = Number of valence electrons + 10 For example, elements of Group 17 have seven valence electrons each. Table 4.4 shows the relationship between the number of valence electrons and the group number.

Table 4.4: Number of Valence Electrons and Group Number

Number of Valence Electrons 1 2 3 4 5 6 7

8 (except

Helium)

Group 1 2 13 14 15 16 17 18

4.3


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Keep in mind that the period number is indicated by the number of filled electron shells. For example, elements in Period 1 each have only one electron shell filled with electrons. All elements in the same period have the same number of filled electron shells. Table 4.5 shows the relationship between the number of shells and the period number.

Table 4.5: Number of Shells and Period Number

Number of Shells Occupied with Electrons

1 2 3 4 5 6 7

Period Number 1 2 3 4 5 6 7

Keep in mind that the chemical properties of the elements in a period are the same but the physical properties of these elements change gradually. For example, the atomic radius of an element (which is half the distance between the nuclei or two atoms of the element joined by a single covalent bond) decreases as it goes across a period from left to right. How about the electronegativity of an element? First, let us remember its definition. Do you still recall? EElectronegativity of an element is the aability of the element to ppull the electron to itself. The electronegativity of elements increases when going across a period from left to right. How about the metal characteristics of the elements? The metal characteristics of the elements decrease when going across the period. It changes from metal to semi-metal and finally to non-metal. The oxide characteristics of the elements change from alkaline to acidic when going across the period.

1. What is the meaning of the term „valence electrons‰?

2. How do we determine the value of valence electrons?

3. What are the periodic patterns noticed in the periodic table?

SELF-CHECK 4.3


1. Fill in the blank of the Periodic Table with the right words to match the appropriate arrows. The first arrow provides an example of metal and non-metal.

Metal� Non�metal� Conductor� Insulator�Ductile� Brittle� Basic�in�water� Acidic�in�water�

Reductants� Oxidants� Cations� Anions�

40 Element X has a proton number of 11.

(a) Draw the electron arrangement of atom X;

(b) State the number of valence electrons of atom X; and

(c) Predict the group and period of atom X in the periodic table.

ACTIVITY 4.2


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TRANSITION ELEMENTS

Firstly, what does it mean by transition elements? TTransition elements are elements in a block located bbetween Group 2 and Group 13 of the periodic table as seen in Figure 4.10. There are 10 elements in each series and they are arranged horizontally.

Figure 4.10: Transition elements of the Periodic Table

We will learn more about these transition elements as we look at their properties and usages in the following subtopics.

4.4.1 Properties

Firstly, keep in mind that all transition elements are mmetals which display the following physical properties:

� Silvery surface;

� Hard;

� High density;

� Ductile and malleable;

� High melting and boiling point; and

� Good electrical conductivity.

4.4


Do you know that transition elements also exhibit four special characteristics which other metals do not have? These special characteristics of transition elements are listed in Figure 4.11.

Figure 4.11: Special characteristics of transition elements

Let us learn more about these four special characteristics: (a) TTransition Elements Form Coloured Compounds Transition elements can form compounds of different colours. Aqueous

solutions of transition element compounds or their ions exhibit certain colours. Table 4.6 shows you the colours of some aqueous solutions of ions of the transition elements.

Table 4.6: Colours of Aqueous Solutions of Ions

Ions CColour

Fe2+ Light green

Fe3+ Brown

Cu2+ Blue

Co2+ Pink

Ni2+ Green

Cr3+ Green

Mn2+ Pink

Cu2+ Blue

MnO4� Purple

CrO42� Yellow

Cr2O72� Orange


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Do you know that the colours of gemstones and precious stones occur naturally due to the presence of the transition elements in them? For example, emerald is green because it contains the transition elements, nickel and iron. Ruby is red due to the existence of chromium. Also, note that aqueous solutions of transition element compounds can react with sodium hydroxide solution and ammonia solution to form coloured precipitates of metal hydroxides. These precipitates may be soluble or insoluble in excess sodium hydroxide solution and ammonia solution. The precipitates formed are coloured because these are compounds of transition elements.

(b) TTransition Elements Have Variable Oxidation Numbers Firstly, what is the function of oxidation number? Oxidation number

measures the charge carried by an element in its compounds. Transition elements show different oxidation numbers in their compounds. This means they can form more than one ion.

Some of the examples are: Iron(II), Fe2+, and Iron(III), Fe3+; and Copper(I),

Cu+, and Copper(II), Cu2+. (c) MMany Transition Elements and Their Compounds Have Catalytic Properties Firstly, what does a catalyst stand for? A ccatalyst is a substance that sspeeds

up a reaction but ddoes not change chemically after a reaction. Many catalysts are transition elements or their compounds. For example:

(i) NNickel

Nickel is used as a catalyst in the manufacture of margarine.

2 2 2 3 3NiCH CH H CH CH� � ��

(ii)� Iron� Iron is used as a catalyst in the Haber process for the manufacture of

ammonia.

2( ) 2( ) 3( )3 2Feg g gN H NH� ��


(d) TTransition Elements Can Form Complex Ions Lastly, transition elements can form complex ions. What is a complex ion?

A ccomplex ion is a ppolyatomic cation or anion consisting of a central metal ion with other groups bonded to it. �An example of a complex ion is tetraamminecopper(II) ion ([Cu (NH3)4]2+) , which consists of four ammonia molecules bonded to the central copper(II) ion.

� � 223 3 44Cu NH Cu NH

��

4.4.2 Industrial Uses

Now, let us look at how these transition elements are used in industrial production. Metals in the transition elements are used widely in society. Some of their usages are shown in Figure 4.12.

Figure 4.12: Industrial uses of transition elements

What are the FOUR characteristics exclusively attributed to transition elements?

SELF-CHECK 4.4


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GROUP 1

Now, let us focus on Group 1. Do you know that Group 1 elements are also known as the alkali metals? What are the elements inside Group 1? The elements of Group 1 are lithium, sodium, potassium, rubidium, caesium and francium (see Table 4.7). These elements are mmetals which can react with water to form alkaline solutions. �

Table 4.7: Elements of Group 1 and their Symbols

Element Element Symbol

Lithium Li

Sodium Na

Potassium K

Rubidium Rb

Caesium Cs

Francium Fr

The following subtopics will explain to you the electronic structure, group trends, and physical and chemical properties of Group 1 elements.

4.5.1 Electronic Structure

Keep in mind that all elements of Group 1 have one valence electron. The atoms of Group 1 elements are able to achieve the stability of a duplex or octet by giving away their valence electron and form singly charged positive ions. You can refer to Table 4.8, which shows the electronic structures of elements in Group 1.

Table 4.8: Electronic Structure of Group 1 Elements

Element Electron Arrangement

Lithium 2.1

Sodium 2.8.1

Potassium 2.8.8.1

Rubidium 2.8.18.8.1

Caesium 2.8.18.18.8.1

Francium 2.8.18.32.18.8.1

4.5


4.5.2 Group Trends

Let us now examine Group 1 trends. Group 1 elements show certain trends when descending the group. For example, the atomic radius increases down the group. This is due to the fact that the number of filled electron shells increases down the group; therefore, the distance between the outermost electron shell and the nucleus increases. The density increases down the group as the increase in mass of the atom is greater than the increase in the atomic radius of the atom. The melting point decreases as we go down Group 1 because the metallic bond between the atoms becomes weaker down the group as the atomic radius increases. How about electropositivity? First, let us know the meaning of electropositivity. Electropositivity is a measurement of the aability of an atom to llose an electron and form a positive ion. As we go down Group 1, the electropositivity of the metal increases because the further the position of the valence electron from the nucleus, the weaker the force of attraction between the nucleus and the valence electron. Hence, the elements lose the single valence electron more easily down the group. As a conclusion for Group 1 trends, you can refer to Table 4.9. It shows some of the properties of Group 1 elements and their trends.

Table 4.9: Properties of Group 1 Elements

Element Proton

No. Nucleon

No. Density (gcm�3)

Hardness (Brinell)

Melting Point ( C)

Boiling Point ( C)

Atomic Radius (nm)

Electro-negativity

Lithium 3 7 0.53 0.06 181 1347 0.15 1.0

Sodium 11 23 0.97 0.07 98 886 0.19 0.9

Potassium 19 39 0.86 0.04 64 774 0.23 0.8

Rubidium 37 85 1.53 0.03 39 688 0.25 0.8

Caesium 55 133 1.87 0.02 28 678 0.26 0.7

Francium 87 223 2.40 n/a 27 677 0.29 0.7


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1. Refer to Table 4.9 and answer the questions based on the explanation given.

State the properties of the elements as you go down the group:

(a) The atomic size;

(b) Hardness;

(c) Melting point;

(d) Boiling point; and

(e) The densities.

2. The above diagram shows the electron arrangement of

Rubidium (Rb) and Caesium (Cs). The further the valence electron is from the nucleus of an atom, the easier it is to remove.

(a) What do you notice about the distance of the valence electron from the nucleus as you go down Group 1?

(b) Why is it so? Relate this to the reactivity of alkali metal.

ACTIVITY 4.3


4.5.3 Physical Properties

What can we say about the physical properties of Group 1 elements? All elements of Group 1 (alkali metals) are ssoft solids and ccan be easily cut. When you cut the alkali metals, you will notice that they resemble grey solids with shiny silvery surfaces. However, their surfaces will turn dull very fast when exposed to air. Alkali metals are very rreactive; hence, they react rapidly with oxygen and water vapour in the air when exposed. Thus, they have to be kept in paraffin oil. Alkali metals have llow densities compared to heavy metals such as iron and copper. They are also good conductors of heat and electricity.

4.5.4 Chemical Properties

Lastly, let us look at the chemical properties of Group 1. As explained earlier, Group 1 elements are also called aalkali metals due to their chemical properties whereby they can readily dissolve in water to form hydroxides, which are strongly alkaline in nature. Keep in mind that they can also form alkaline oxides. Group 1 elements exhibit similar chemical properties in their reactions with the following elements:

(a) React with water to liberate hydrogen gas and form metal hydroxide;

(b) React with oxygen to produce metal oxides;

(c) React with chlorine to produce metal chloride; and

(d) React with bromine to produce metal bromide.

1. Watch this video to see the reactions of Group 1 metals to water at http://www.youtube.com/watch?v=Ft4E1eCUItI&feature=player_embedded. After watching the simulation, what can you conclude?

2. Find out how you can investigate the chemical properties of Group 1 alkali metals when they react with oxygen.

3. How do we investigate the chemical properties of Group 1 metals in their reactions with chlorine and bromine?

ACTIVITY 4.4


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GROUP 17

Let us learn more about Group 17. What are the elements inside Group 17? Group 17 elements are fluorine, chlorine, bromine, iodine and astatine (see Table 4.10).

Table 4.10: Elements of Group 17 and Element Symbols

Element EElement Symbol MMolecular Formulae

Fluorine F F2

Chlorine CI Cl2

Bromine Br Br2

Iodine I I2

Astatine At At2

Do you know that they are known as „„halogens‰? This comes from a Greek word for „salt producer‰. Halogens are vvery reactive elements and most of them exist naturally as hhalide salts. Halogens are nnon-metals and exist as ddiatomic covalent molecules. Next, let us investigate the electronic structure, group trends, physical properties, and chemical properties of Group 17.

4.6.1 Electronic Structure

Firstly, note that all elements of Group 17 have seven valence electrons. The atoms of Group 17 elements are able to achieve the stability of an octet arrangement by accepting one electron and form negative ions. Table 4.11 shows the electronic structures of elements in Group 17.

4.6

SELF-CHECK 4.5

1. How many valence electrons does an atom of rubidium have?

2. Compare the melting point of rubidium and potassium. Explain your answer.

3. A small piece of rubidium is added to cold water. Observe what happens and write an equation for the reaction.


Table 4.11: Electronic Structure of Group 17 Elements


Flourine 2.7

Chlorine 2.8.7

Bromine 2.8.18.7

Iodin 2.8.18.18.7

Astatine 2.8.18.32.18.7

4.6.2 Group Trends

How about the group trends of Group 17 elements? Group 17 elements show a certain trend when descending the group. For example, the atomic radius increases gradually down the group. This is because the number of shells occupied with electrons increases down the group, resulting in an increase in the distance between the outermost electron shell and the nucleus. How about the density? The density increases down the group due to the fact that the increase in atomic mass is greater than the increase in volume down the group. Generally, the halogens have low melting and boiling points because the forces of attraction between the molecules are weak. However, the melting and boiling points of halogens increase down the group because the molecular size increases down the group. As the molecular size of halogens increases, the van der Waals forces of attraction between the molecules become stronger. Therefore, more heat is required to overcome the forces of attraction between the molecules, resulting in an increase in the melting and boiling points. For instance, the first two elements (fluorine and chlorine) have low boiling points and are gases at room temperature. Bromine is a liquid whereas iodine and astatine are solids at room temperature.


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How about their colours? The colour of halogen becomes darker down the group. For example, fluorine is a pale-yellow gas, chlorine is a greenish-yellow gas, and bromine is a reddish-brown liquid whereas iodine is a purplish-black solid. Lastly, let us look at the trend of electronegativity. Firstly, what does it stand for? EElectronegativity is a measurement of the tendency of an atom to aattract an electron and form a negative ion. Generally, all halogens are eelectronegative non-metals. However, the electronegativity of halogens decreases from fluorine to iodine. This is because as the atomic radius becomes larger down the group, the force of attraction between the nucleus and the valence electrons becomes weaker; hence, the strength of the nucleus to attract electrons becomes weaker. Table 4.12 shows some of the properties of Group 17 elements and their trends.

Table 4.12: Properties of Group 17 Elements and their Trends

Element Fluorine Chlorine Bromine Iodine

Proton number 9 17 35 53

Atomic radius (nm) 0.071 0.099 0.114 0.133

Density (gcm�3) 0.0017 0.0032 3.13 4.94

Melting point ( C) �220 �101 �7 114

Boiling point ( C) �188 �35 59 184

Colour Pale-yellow Greenish- yellow

Reddish-brown

Purplish-black

Electronegativity 4.0 3.0 2.8 2.5

Source: http://www.rsc.org


4.6.3 Physical Properties

How about the physical properties of Group 17? Earlier, we learnt that all Group 17 elements are non-metals. Therefore, they are iinsulators of heat and electricity. Generally, all halogens have llow melting and boiling points due to poor forces of attraction between the molecules. All halogens have llow densities.

4.6.4 Chemical Properties

Lastly, let us check out the chemical properties of Group 17. Keep in mind that all Group 17 elements have seven valence electrons and react in a similar manner due to the electron arrangement. The reactivity of halogen decreases as they go down the group. Halogens can easily gain one electron to achieve a stable octet electron arrangement and therefore are good oxidising agents. As the reactivity of halogens decreases down the group, their strength as oxidising agents also decreases down the group. Note that Group 17 elements exhibit similar chemical properties when they react with:

(a) Water;

(b) Iron; and

(c) Cold sodium hydroxide solution. Halogens react with water to produce acidic solutions. However, the solubility of halogens in water decreases when going down group 17. For example, chlorine and bromine dissolve readily in water, forming acidic solutions which turn blue litmus paper red. The solution formed is also bleaching agents which turn the litmus paper white due to the presence of hypochlorous acid or hypobromus acid. On the other hand, iodine is only very slightly soluble in water. Chlorine, bromine and iodine react with hot iron to produce a brown iron(III) halides solid. The reactivity of the halogens in their reaction with iron decreases from chlorine to bromine. The halogens react with cold sodium hydroxide solution to produce water and a colourless solution containing salts of sodium halide and sodium halate(I). The reactivity of the halogens in their reaction with cold sodium hyroxide solution decreases from chlorine to iodine.


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NOBLE GASES

Now we move on to noble gases. Generally, what can we say about noble gases? Noble gases are elements of GGroup 18, consisting of helium, neon, argon, krypton, xenon and radon, as shown in Table 4.13.

Table 4.13: Elements of Noble Gases (Group 18) and their Symbols

Element Element Symbol

Helium He

Neon Ne

Argon Ar

Krypton Kr

Xenon Xe

Radon Rn

Keep in mind that noble gases are chemically unreactive because the atoms have a stable electronic structure. Helium atoms have two electrons in their only electron shell and this electronic structure is also known as the dduplet electron arrangement. On the other hand, the outermost shell of the atoms of the other noble gases has eight electrons and this electronic structure is known as the ooctet electron arrangement. Table 4.14 shows the electronic structure of the atoms of noble gases.

Table 4.14: Electronic Structure of Noble Gases Elements


Helium 2

Neon 2.8

Argon 2.8.8

Krypton 2.8.18.8

Xenon 2.8.18.18.8

Radon 2.8.18.32.18.8

4.7


Generally, all noble gases have these characteristics:

(a) Exist as monoatom gases;

(b) Colourless and insoluble in water;

(c) Cannot conduct electricity and are poor conductor of heat;

(d) Have very low melting and boiling points because noble gases are held together by weak Van der Waals forces of attraction; and

(e) Have very low densities but the density increases slowly when going down the group because the increase in atomic mass is greater than the increase in volume.

Lastly, do you know that noble gases are used in various activities and equipment in our daily lives? These are due to their inert property. Table 4.15 shows some physical properties of noble gases elements.

Table 4.15: Properties of Noble Gases Elements

Element Helium Neon Argon Krypton Xenon Radon

Proton number 2 10 18 36 54 86

Atomic radius (nm)

0.050 0.070 0.094 0.109 0.130 �

Density (gcm�3) 0.17 0.84 1.66 3.45 5.45 �

Melting point ( C) �270 �248 �189 �156 �112 �71

Boiling point ( C) �269 �246 �186 �152 �107 �62

Source: http://www.rsc.org

�

ACTIVITY 4.5

Can you find the usage of noble gases in various activities and equipment in our daily lives? You can search the Web and have discussions with your friends during tutorial class. Good luck!


103

PERIOD 3

Before we begin this lesson, let us recall the meaning of „period‰. As stated before, period refers to the elements in each horizontal row of the periodic table. How about Period 3? Period 3 elements are located in the third row of the periodic table. This row consists of elements such as sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine and argon (see Table 4.16).

Table 4.16: Period 3 Elements and the Element Symbol

Element EElement Symbol

Sodium Na

Magnesium Mg

Aluminium Al

Silicon Si

Phosphorus P

Sulphur S

Chlorine CI

Argon Ar

What are the properties for the elements in Period 3? They are explained further in Table 4.17. �

Table 4.17: Properties of the Elements in Period 3

Element Na Mg Al Si P S Cl Ar

Proton Number 11 12 13 14 15 16 17 18

Electron arrangement

2.8.1. 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8

Atomic Radius (nm)

186 160 143 118 110 104 100 94

Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 �

Melting Point ( C) 98 649 660 111 44 113 �101 �189

Boiling Point ( C) 886 1090 2467 2355 280 444 �35 �186

Source: http://www.chemguide.co.uk

4.8


Do you know that elements of Period 3 can be classified into metals and non-metals? This is based on the basic or acidic properties of their oxides which can determine the metallic or non-metallic properties of the elements. MMetals form oxides with bbasic properties only. Some mmetals can form oxides with bboth acidic and basic properties. These oxides are known as aamphoteric oxides. NNon-metals form oxides with aacidic properties only. Properties of oxides of the elements in Period 3 can be summarised as in Table 4.18.

Table 4.18: Properties of Oxides of the Elements in Period 3

Oxides of the Elements in Period 3

Na2O MgO Al2O3 SiO2 P4O10 SO2 Cl2O7

Properties of Oxide Basic Basic Amphoteric Acidic Acidic Acidic Acidic

Based on Table 4.18, we can conclude that properties of oxides of the elements change from basic to acidic when going across Period 3 from left to right.

4.8.1 Chloride and Hydride for Elements in Period 3

Before we end this topic, let us learn more about chloride and hydride for Period 3 elements. You can refer to Table 4.19, which shows chloride and hydride for elements in Period 3.

Table 4.19: Chloride and Hydride for Elements in Period 3

Group 1 2 3 4 5 6 7

Formula of chloride

NaCl MgCl2

AlCl3 SiCl4 PCl5 S2Cl2 Cl2

Melting point (K)

1074 987 463 203 435 193 172

Bonding Ionic Ionic Covalent Covalent Covalent Covalent Covalent

Solubility in water

Very Very Hydrolysed Hydrolysed Hydrolysed � �

pH of solution

7 7 3 0 0 � �


105

Based on Table 4.19, we can deduce the following when going across Period 3 from left to right:

(a) The bonding changes from ionic to covalent;

(b) They react with water rather than dissolve in it; and

(c) The compound solution changes from neutral to acidic.

1. The letters A to K are used to represent a few elements in the periodic table as shown below.

(a) With reference to the above table, name the element that is:

(i) A halogen;

(ii) A transition metal;

(iii) An inert gas; and

(iv) The most reactive metal.

(b) Why is inert gas inactive?

(c) State the element which forms coloured ions in an aqueous solution.

(d) Element K has a proton number of 23 and nucleon number of 51. How many neutrons and electrons are there in a K5+

ion? Write the oxide formula of the K5+ ion.

2. The aqueous solutions of ions of transition elements have certain colours. What are the colours of the following solutions?

(a) Potassium Manganate(VII)

(b) Iron(II) Sulphate

ACTIVITY 4.6


� The elements in the periodic table are arranged in rows of periods and

columns called the groups.

� The development of the periodic table began with Antoine LavoisierÊs effort in classifying elements into groups. Over the years, improvements were made to the table, resulting in the birth of the modern day periodic table revised by Glenn Seaborg.

� The number of valence electrons in an atom decides the position of the group of that element in the periodic table.

� Some groups have special names:

� Group 1 elements are called alkali metals;

� Group 2 elements are known as alkaline earth metals;

� Group 3 to Group 12 elements are known as transition elements;

� Group 17 elements are called halogens; and

� Group 18 elements are called noble gases.

� Elements are arranged horizontally in the ascending order of their proton numbers in the periodic table.

3. Name the transition element which is used as a catalyst in the following processes.

(a) Contact process.

(b) Haber process. 4. Cu2+ ions react with ammonia solution to form sediment.

(a) State the colour of the sediment that forms initially.

(b) Write the above equation of the reaction.

(c) Why does the sediment become soluble again when excess ammonia solution is added?


107

� Elements in a group have very similar chemical properties.

� The atomic radius increases when going down the group and decreases when going across a period from left to right.

� The electronegativity of an element is the ability of the element to pull the electron towards itself. The electronegativity of elements increases when going across a period from left to right.

� The metal characteristic of the elements decreases when going across the period. It changes from metal to semi-metal and finally to non-metal.

� The oxide characteristics of elements change from alkaline to acidic when going across the period.

� The outer electron is also known as the valence electron. It is found in the outermost occupied shell of an atom.

� Transition elements are elements in a block located between Group 2 and Group 13 of the periodic table.

� Aqueous solutions of transition element compounds can react with sodium hydroxide solution and ammonia solution to form coloured precipitates of metal hydroxides.

� Transition elements are metals that have four special characteristics which other metals do not have:

� They form coloured compounds;

� They have variable oxidation numbers;

� They have catalytic properties; and

� They form complex ions.

� Metals in the transition elements have industrial uses and are used in the production of paints, glass, vegetable oil, sulphuric acid, nitric acid and ammonia.

� Group 1 elements are also called alkali metals because they readily dissolve in water to form hydroxides, which are strongly alkaline in nature. They also form alkaline oxides.

� All alkali metals are soft grey solids with shiny silvery surfaces when freshly cut.


� Alkali metals are very reactive and they react rapidly with oxygen and water vapour in the air when exposed.

� Alkali metals have low densities compared to heavy metals such as iron and copper.

� They are also good conductors of heat and electricity.

� Group 1 elements exhibit similar chemical properties in their reactions with:

� Water to liberate hydrogen gas and form metal hydroxide;

� Oxygen to produce metal oxides;

� Chlorine to produce metal chloride; and

� Bromine to produce metal bromide.

� Group 17 elements are fluorine, chlorine, bromine, iodine and astatine.

� They are known as halogens and are reactive non-metals. They exist as diatomic covalent molecules.

� Group 17 elements exhibit similar chemical properties in their reactions with:

� Water to produce two types of acids;

� Iron to produce iron(III) halides; and

� Sodium hydroxide solution to produce two types of sodium salts and water.

� Elements of Period 3 can be classified as metals and non-metals based on the basic or acidic properties of their oxides.

� Metals form oxides with basic properties only. Some metals can form oxides with both acidic and basic properties. These oxides are known as amphoteric oxides. Non-metals form oxides with acidic properties only.

� When going across Period 3 from left to right, the oxide properties of elements change from basic to acidic.

� When going across Period 3 from left to right:

� The bonding changes from ionic to covalent;

� They react with water rather than dissolve in it; and

� The compound solution changes from neutral to acidic.


109

Actinides

Alkali metals

Alkaline earth metals

Atomic radius

Electronegativity

Groups

Halogens

Lanthanides

Long periods

Metals

Noble gases

Non-metals

Oxide characteristics

Periodic table

Periods

Short periods

Transition elements

Valence electrons

Briggs, J. G. R. (2003). Science in focus chemistry for GCE ÂOÊ level. Singapore:

Pearson Education Asia Pte Ltd.

Conoley, C., & Hills, P. (2002). Chemistry (2nd ed.). London: Harper-Collins.

Hewitt, P. G. (1998). Conceptual physics (8th ed.). Massachusetts: Addison-Wesley.

Kementerian Pendidikan Malaysia � Bahagian Pendidikan Guru. (1995). Buku sumber pengajaran pembelajaran sains sekolah rendah: Strategi pengajaran dan pembelajaran sains. Kuala Lumpur: Kementerian Pendidikan Malaysia.

Ralph, A. B. (2003). Fundamentals of chemistry. New Jersey: Prentice Hall.

Whitten, K. W., Davis, R. E., Peck, M. L., & Stanley, G. G. (2010). Chemistry (9th ed.). Belmont: Brooks/Cole.

topic 4 periodic table

Documents

modern periodic table

classification table

discovered elements

usages of transition

similar properties

group trends

antoine lavoisier figure

classification of period