Sec 2 Term 1_1 Notes

8/15/2019 Sec 2 Term 1_1 Notes

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Term 1_1 Notes: Periodic Table, Atomic Structure

Name: _______________________ ( ) Class: 2 ___ Date: _________

Atomic s tru ctu re

What is an atom?

An atom is the smallest unit of an element, having the properties of that element.

Are there particles smaller than the atom?

Atoms are not like solid balls (Figure 1) as proposed by Dalton in 1803.


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The key points about atomic structure are summarised below:

Thinking Time!

Why is the term "relative mass" used in the Table 1 rather than just mass?


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Advanced: For elements after calcium in the 4th period, their third shell can hold up to 18electrons.

Some interesting facts...

• The nucleus takes up less than 1% of the volume of a nucleus.• More than 99% of an atom is empty space occupied by rapid moving electrons.

Protons, neutrons and electrons are the building blocks for all atoms. Hydrogen atoms are thesimplest atoms. Each hydrogen atom has only one proton and one electron. The next

simplest atoms are those of helium with two protons, two electrons and two neutrons. Afterhelium comes lithium, with three protons, three electrons and four neutrons.


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Atomic number and mass number

The only atoms with one proton are those of hydrogen. The only atoms with two protons arethose of helium. The only atoms with three protons are those of lithium and so on.

This shows that the number of protons in an atom decides which element it is. Because ofthis, the number of protons in an atom is called its atomic number or proton number. Thus,hydrogen has an atomic number of one, helium has an atomic number of two, lithium three,and so on. Remember also that the order of the element in the periodic table tells you itsatomic number. So, chlorine, the seventeenth element in the periodic table with 17 protons

and 17 electrons has an atomic number of 17.

The mass of the electrons in an atom is negligible compared to that of the protons andneutrons. In fact, the mass of an atom depends on the number of protons and neutrons addedtogether. This number is called the mass number or nucleon number of the atom.

So,

atomic number = number of protons

mass number = number of protons +

number of neutrons

So, aluminium atoms, with 13 protons and 14 neutrons, have an atomic number of 13 and a massnumber of 27. Sometimes the symbol Z is used for atomic number and the symbol A for mass

b S f l i i Z 13 d A 27


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Thinking Time!

An atom can be described as an electrically neutral entity made up of a positively chargednucleus at its centre with negatively charged electrons moving around the nucleus.

(a) Why is the atom electrically neutral?

(b) Why is the nucleus positively charged?

Isotopes and Relative Atomic Mass

Several elements have relative atomic masses which are whole numbers. For example, therelative atomic mass of carbon is 12.0, that of fluorine is 19.0 and that of sodium is 23.0. Thisis not surprising, as the mass of an atom depends on the mass of its protons and neutrons, both ofwhich have a relative mass of 1.0. For example, we could calculate the relative mass of

fl i f ll


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Each isotope has a relative atomic mass which is a whole number, but the average relative atomicmass for the mixture of isotopes is not always a whole number.

Chlorine is a good example of an element with isotopes. Naturally occurring chlorine contains twoisotopes 3517

C1 called chlorine-35 and 3717 C1 called chlorine-37. Each of these isotopes has 17 protons

and 17 electrons. Therefore, both isotopes have the same atomic number and the same chemical properties because these are determined by the number of electrons.

However, one isotope (3517 C1 ) has 18 neutrons and the other (3717

C1) has 20 neutrons. Therefore,

they have different mass numbers, different masses and hencedifferent physical properties

becausethese depend on the masses of atoms and molecules.

The similarities and differences between isotopes of the same element are summarised in the table below.

The similarities and differences between isotopes of the same element

I t b di id d i t t t O t i di ti th th i di ti


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Radioactive isotopes produce heat which can beused to produce electrical energy. Small amounts

of radioactive isotopes are used to supply energyin remote places. Radioactive isotopes provideenergy for spacecraft exploring the outer planets,such as Jupiter and Saturn.

Some people have irregular heartbeats. They needto have a heart pacemaker implanted inside theirchest. This instrument provides a tiny electricalshock to ensure a steady heartbeat. Pacemakerscan be powered by radioactive isotopes. Such a

pacemaker can work reliably for over 20 years. Anordinary battery would have to be replaced every10 years.

Nuclear Energy (Optional)

Radioactive isotopes are used to produce large amounts of energy iinuclear fission. This is done in acontrolled way in a nuclei reactor.The fuel is an isotope of uranium,

i 235 I th t th


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Calculating relative atomic masses (Optional)

The relative atomic mass of an element is the average mass of one atom. This can becalculated from the relative masses of its isotopes and their relative proportions.

Look closely at the figure below which shows a mass spectrometer trace for chlorine. The traceshows that chlorine contains two isotopes, with mass numbers of 35 and 37.


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Electronic Structure & Electronic Configuration

Nitrogen-14 atom has 7 electrons. Two of its electrons will go into the 1st shell; theremaining electrons will go into the 2nd shell (Figure 2). With 7 electrons, nitrogen has theelectronic configuration of 2.5

Argon-40 atom has 18 electrons. Two of its electrons will go into the 1st shell, 8 electronswill go into the 2nd shell, and the remaining 8 electrons will go into the 3rd shell (Figure 3).With 18 electrons, argon has the electronic configuration of 2.8.8

Figure 2: Electronic Structure (FullElectronic Configuration) of Nitrogen-14atom

Figure 3: Electronic Structure (FullElectronic Configuration) of Argon-40atom

(b) Valence Shell (Outer Shell)


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(a) Formation of cations

When an atom loses one or more electrons, it becomes a positively charged particle calledcation.

Lithium atom (Li)3 electrons3 protons

Net charge: 0

Lithium ion (Li+)2 electrons3 protons

Net charge: +1

In a lithium atom, there are 3 protons and 3 electrons. In a lithium ion, there are 3 protonsand 2 electrons. Therefore, the lithium ion carries an overall positive charge of 1+ and iswritten as Li+.

(b) Formation of anions

When an atom gains one or more electrons, it becomes a negatively charged particle calledanion.


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The Periodic Table

During the 19th century, several chemists looked for patterns in the properties of elements. Themost successful of these approaches was by the Russian chemist Dmitri Mendeleev in 1869.

• Mendeleev arranged all the known elements in order of their relative atomic masses.

• He also arranged the elements in horizontal rows so that elements with similar propertieswere in the same vertical column.

Because of the periodic repetition of elements with similar properties, Mendeleev called his

arrangement a periodic table.

The figure below shows part of Mendeleev's periodic table. Notice that elements with similar

properties, such as sodium and potassium, fall in the same vertical column. Which other pairs or

trios of similar elements appear in the same vertical column of Mendeleev’s table?

In the periodic table

• The vertical

columns

of similar elements

are called groups.

• The horizontal rows


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Period - a horizontal row of elements

(a) Group


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Element Proton

number

Number of electrons in Electronic

configuration

Period Group

1st shell

2n shell

3r shell

4t shell

H 1 1 1 1 -

He 2 2 2 1 0

Li 3 2 1 2.1 2 I

Be 4 2 2 2.2 2 II

B 5 2 3 2.3 2 III

C 6 2 4 2.4 2 IV N 7 2 5 2.5 2 V

O 8 2 6 2.6 2 VI

F 9 2 7 2.7 2 VII

Ne 10 2 8 2.8 2 0

Na 11 2 8 1 2.8.1 3 I

Mg 12 2 8 2 2.8.2 3 II

Al 13 2 8 3 2.8.3 3 III

Si 14 2 8 4 2.8.4 3 IV

P 15 2 8 5 2.8.5 3 V

S 16 2 8 6 2.8.6 3 VI

Cl 17 2 8 7 2.8.7 3 VII

Ar 18 2 8 8 2.8.8 3 0

K 19 2 8 8 1 2.8.8.1 4 I

Ca 20 2 8 8 2 2.8.8.2 4 II


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They conduct electricity better than non-metals but not as well as metals. Elements with some

properties like metals and other properties like non-metals are called metalloids. Because of this

difficulty in classifying elements neatly as metals and non-metals, chemists looked for patternsin the properties and reactions of smaller groups of elements.

The modern periodic table

The modern periodic table, shown over the next page, is based on Mendeleev's. It shows all theknown elements numbered along each period, starting with period 1, then period 2, etc. The number

given to each element is called its atomic number. Thus, hydrogen has an atomic number of 1,helium 2, lithium 3, etc.

The periodic table shows all the elements arranged in order of increasing atomic number.

(i) Electronic Structure

Down each group, the number of valence electrons is the same for each element and is equalto the group number.

Example: Group I Elements

Element Electronic configuration

Li 2.1

N 2 8 1


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(ii) Metals and non-metals

Across the period, the properties of elements change from metallic to non-metallic.

Generally, elements with small number of electrons in the valence shell (e.g. Group I and II)are metals. Elements with large number of electrons in the valence shell (e.g. Group VII and0) are non-metals.

The line that divides metals from non-metals runs run diagonally through the Periodic Table.Elements found beside this dividing line (zigzag line) are known as metalloids. Metalloidshave some properties of non-metals and metals.


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Electron structure and chemical bonding

We have previously looked at the first twenty elements in the periodic table.

When elements react we now know that they try to gain, lose or share electrons in order

to get a more stable electron structure. In many cases, this more stable electron

structure is the same as that of a noble gas.

When elements react we now know that they try to gain, lose or share electrons in order toget a more stable electron structure. In many cases, this more stable electron structure is thesame as that of a noble gas.

The simple ideas expressed in this statement form the basis of the electronic theory ofchemical bonding.

Look carefully at the table below. This shows the electron structures of the atoms and ions ofelements in period 3.


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Why do atoms bond?

Noble gases, such as helium, neon and argon, are monoatomic because their valence shellsare fully occupied by electrons. Thus, noble gases are stable and do not undergo bondingwith other atoms.

Since atoms with electronic configurations of noble gases are stable, atoms bond to achieveelectronic configuration of a noble gas. Atoms do so by transfer or sharing of electronswith other atoms.

By having an electronic configuration of a noble gas, an atom will achieve stability.

- When you have 2 electrons in the 1st shell, you have a duplet structure.- When you have 8 electrons in the rest of the shells, you have an octet structure.

Chemical bonds

There are three ways of forming chemical bonds between atoms:(i) Ionic bonding(ii) Covalent bonding(iii) Metallic bonding (to be covered in Sec. 3)

i. Ionic bonding


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In a chlorine atom, there are 17 protons and 17 electrons. In a chloride ion, there are 17 protons and 18 electrons. Therefore, the chloride ion carries an overall positive charge of 1-

and is written as Cl-

.

Figure 5: Formation of Sodium Chloride

In figure 6, the magnesium atom gives up two electrons to form a magnesium ion, Mg 2+.These two electrons are transferred to two chlorine atoms to form two chloride ions, Cl-. Themagnesium chloride has the formula MgCl2.


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(b) ‘Dot-and-cross’ diagrams of ionic compounds

Steps:1. Draw the valence electrons of the elements using ‘Dot and Cross’ only.2. Electron(s) is/are transferred from the valence shell of the metal to the valence shell of

the non-metal.3. The anion that has gained electron(s) from the cation will now have two type of

electrons – one originally from its valence shell, another one from the cation –differentiated by ‘dots’ and ‘crosses’.

4. Remember to indicate the charges of the ions.

Examples of ‘Dot-and-cross’ diagram can be found in figure 7 and 8.

Figure 7: ‘Dot-and-cross’ diagram showing the bonding in sodium chloride


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(a) Covalent bonds in elements

Hydrogen molecules, H2

A hydrogen atom has 1 valence electron. It requires one more electron to obtain theelectronic configuration of a noble gas.

In order to obtain the electronic configuration of helium, two hydrogen atoms can share a pairof electrons between themselves to form hydrogen molecules, H2 (figure 5). In order todifferentiate the two electrons, the electron of one atom is represented by a ‘cross’ while theother electron of another atom is represented by a ‘dot’.

Formation of hydrogen molecule

Oxygen molecules, O2

In order to obtain the electronic configuration of a noble gas, each oxygen atom requires twomore electrons. Instead of sharing a pair of electrons, two oxygen atoms can share two pairsof electrons to form a double bond and obtain an octet structure.


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(b) Covalent bonds in compounds

Methane molecule, CH4

The carbon atom has four valence electrons and it needs four more electrons to gain an octetstructure. The carbon atom can share its four electrons with four other hydrogen atoms,forming a single covalent bond with each of the hydrogen atom.

Formation of methane molecule (showing of electrons in the outer shells only)


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Structure and Properties

Look at the crystals of sodium chloride in this photograph. What do you notice about all the saltcrystals?

All the salt crystals are roughly the same cubicshape. Further studies show that all the crystals ofone substance have similar shapes. This suggests

that the particles in the crystals are always packedin a regular fashion to give the same overall shape.Sometimes, crystals grow unevenly and theirshapes become distorted. Even so, it is usually easyto see their general shape. Solid substances whichhave a regular packing of particles are described ascrystalline. The particles may be atoms, ions ormolecules.

The figure below shows how cubic crystals and hexagonal crystals can form. If the particlesare always placed in parallel lines or at 90o to each other, the crystal will be cubic. If the

particles are placed at 120° in the shape of a hexagon, the final crystal will be hexagonal.


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Using X-rays to study crystals (Optional)

Look through a piece of thin stretched cloth at a small bright light. The pattern you see is dueto the deflection of the light as it passes through the regularly spaced threads of the fabric.This deflection of the light is called diffraction and the patterns produced are diffractionpatterns. If the cloth is stretched so that the threads in the fabric get closer, then the patternspreads further out. From the diffraction pattern which we can see, we can work out the

pattern of the threads in the fabric which we cannot see. The same idea is used to work outhow the particles are arranged in a crystal.

A narrow beam of X-rays is directed at a well-formed crystal. Some of the X-rays are diffracted

by particles in the crystal onto X-ray sensitivefilm. When the film is developed, a regular

pattern of spots appears. This is the diffraction pattern for the crystal. From the diffraction pattern which we can see, it is possible to work

out the pattern of particles in the crystal which wecannot see. A regular arrangement of spots on thefilm indicates a regular arrangement of particlesin the crystal. This regular arrangement of

particles in the crystal is called a lattice.

X-rays have been used in this way to study the structure of thousands of different solids.Beams of electrons can also be used, like X-rays, to study the way in which particles are


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Notice that the structure and bonding of a substance determine its properties and, in

turn, the properties determine its uses.

So the links from structure and bonding to properties help us to explain the uses of materials -why metals are used as conductors, why graphite is used in pencils and why clay is used tomake bricks.

Earlier, you have learnt that all substances are made up from only three different types ofparticle - atoms, ions, and molecules.

These three particles give rise to four different solid structures.

• giant metallic structures,

• giant covalent structures,

• giant ionic structures, and

• simple molecular structures.

The table below shows the particles in these four structures, the types of substances formedand examples of these substances.

The four types of solid structure and the particles they contain

Types of

structure

Particles in the

structure

Types of substance Examples

Giantmetallic

atoms Metals and alloys(mixture of metals)

Na, Fe, Cu, steel, brass


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X-ray analysis shows that the atoms in metal grains - packed in a regular fashion, but thegrains themselves are irregular-shaped crystals pushed tightly together.

Metals usually have a high density. This suggests that the atoms are packed close together. Infact, X-ray studies show that the atoms of most metals are packed as close together as

possible. This arrangement is close packing. The figure above shows a few atoms in layer ofa metal crystal.

Notice that each atom in the middle of thecrystal touches six other atoms in the samelayer. When a second layer is placed on top ofthe first layer, atoms in the second layer sinkinto the dips between atoms in the first layer.This means that any one atom in the first layercan touch six atoms in its own layer, threeatoms in the layer above it and three atoms in

the layer below, i.e. a total of twelve atoms inall.

The properties of metals


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• Good conductivity - When a metal isconnected in a circuit, freely moving electrons

in the metal move towards the positiveterminal. At the same time, electrons move intothe other end of the wire from the negativeterminal. This flow of electrons through thewire forms the electric current.

• Malleability - The bonds between atoms in ametal are strong but they are not rigid. When aforce is applied to a metal crystal, the layers ofatoms can `slide' over each other. This isknown as slip. After slipping, the atoms settleinto position again and the close-packedstructure is restored.

This diagram shows the positionsof atoms before and after slip.This is what happens when ametal is bent or hammered intodifferent shapes.


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diamond is a single giant molecule or a macromolecule. Only a small number of atoms areshown in the model. In a real diamond, there are billions of atoms.

The properties of diamond

• Diamond is very hard because its carbon atoms are linked by very strong covalent bonds. Another reason for its hardness is that the atoms are not arranged in layers so theycannot slide over one another like the atoms in metals. In fact, diamond is the hardestknown natural substance. Most of its industrial uses depend on this hardness.

• Diamond has a very high melting point because of the strong covalent bonds linkingcarbon atoms in a giant structure. This means that the atoms cannot vibrate fast enough to

break away from their neighbours until very high temperatures are reached.

• Diamond does not conduct electricity. Unlike metals, diamond has no free electrons because all four electrons in the outer shell of each carbon atom are held firmly incovalent bonds. So in diamond there are no free electrons to form an electric current.

Giant Ionic Structures


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Simple molecular substances

Oxygen and water are good examples of simplemolecular substances. They are made of simplemolecules each containing a few atoms. Their


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For example, hydrogen is H2, chlorine is Cl2, iodine is I2, carbon dioxide is CO2 andtetrachloromethane is CCl4 . Sugar (C12H22O11) has much larger molecules than these

substances, but it still counts as a simple molecule.

In these simple molecular substances, the atoms are held together in each molecule by strongcovalent bonds. But there are only weak forces between the separate molecules. These weakforces between the separate molecules are called intermolecular bonds or Van der Waalsforces.

The properties of simple molecular substances

The properties of simple molecular substances can be explained in terms of their structure.


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• Simple molecular substances do not conduct electricity. They have no mobileelectrons like metals. They do not have any ions either. This means that they cannot

conduct electricity as solids, as liquids or in aqueous solution.

Notice the following key points from the last four ions.

• Substances with giant structures are often hard with high melting points and boiling points.

• Substances with simple molecular structures are usually soft with low melting pointsand boiling points.

• There are three types of strong force between particles in giant structures; metallic bonds between metal atoms, covalent bonds between non-metal atoms, ionic bonds between positive metal ions and negative non-metal ions.

• In simple molecular substances there are relatively weak forces between the separatemolecules.

Ambook Resources

For more fun and interactivity of this topic, do refer to your amBook under theunit: Interactions – Atoms and Molecules.


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MODERN CHEMICAL SYMBOLS

Listed below are the atomic numbers, names, and symbols of the most common elements. Theatomic number is used to determine the place of the element in the periodic table; it also hasother meaning as you will find out later in the course.

Become familiar with the names and symbols of these elements.

Atomic Atomic

Number Name Symbol Number Name Symbol

———— ——————— ———— ———— ——————— ————

1 hydrogen H 28 nickel Ni

2 helium He 29 copper Cu

3 lithium Li 30 zinc Zn

4 beryllium Be 33 arsenic As

5 boron B 35 bromine Br

6 carbon C 36 krypton Kr

7 nitrogen N 37 rubidium Rb

8 oxygen O 38 strontium Sr

9 fluorine F 47 silver Ag

10 neon Ne 48 cadmium Cd

11 sodium Na 50 tin Sn

12 magnesium Mg 51 antimony Sb

13 aluminum Al 53 iodine 1


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Common ions and formulae of ionic compounds

Symbols of common ionsYou are not expected to know all the names and symbols of common ions, but you should beable to work out the formulae of ionic compounds. The names and symbols of some ions areshown below.

Formulae of ionic compounds

Ionic compounds contain positive and negative ions. The number of positive charges must


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Watch out for compound ions, e.g. ammonium, hydrogencarbonate, hydroxide, nitrate,sulfate and carbonate. If you need more than one of them to balance the charges, put brackets

around their symbol at step (a) or (c). For example, sodium hydroxide is NaOH, butmagnesium hydroxide is Mg(OH)2; copper (II) sulfate is CuSO4, but ammonium sulfate is(NH4)2SO4.


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35 Notes compiled by Lower Sec Science department @HCI

Group

I II III IV V VI VII 01 H

Hydrogen

1

4

He Helium

2 7

Li Lithium

3

9

Be Beryllium

4

11

B Boron

5

12

C Carbon

6

14

N Nitrogen

7

16

O Oxygen

8

19

F Fluorine

9

20

Ne Neon

10 23

Na Sodium

11

24

Mg Magnesium

12

27

Al Aluminium

13

28

Si Silicon

14

31

P Phosphorus

15

32

S Sulphur

16

35.5

Cl Chlorine

17

40

Ar Argon

18 39

K Potassium

19

40

Ca Calcium

20

45

Sc Scandium

21

48

TiTitanium

22

51

V Vanadium

23

52

Cr Chromium

24

55

Mn Manganese

25

56

Fe Iron

26

59

Co Cobalt

27

59

Ni Nickel

28

64

Cu Copper

29

65

Zn Zinc

30

70

Ga Gallium

31

73

Ge Germanium

32

75

As Arsenic

33

79

Se Selenium

34

80

Br Bromine

35

84

Kr Krypton

36

85

Rb Rubidium

37

88

Sr Strontium

38

89

Y Yttrium

39

91

Zr Zirconium

40

93

Nb Niobium

41

96

Mo Molybdenum

42 Tc

Technetium

43

101

Ru Ruthenium

44

103

Rh Rhodium

45

106

Pd Palladium

46

108

Ag Silver

47

112

Cd Cadmium

48

115

In Indium

49

119

Sn Tin

50

122

Sb Antimony

51

128

Te Tellurium

52

127

I Iodine

53

131

Xe Xenon

54 133

Cs Caesium

55

137

Ba Barium

56

139

La Lanthanium

57

178

Hf Hafnium

72

181

Ta Tantalum

73

184

W Tungsten

74

186

Re Rhenium

75

190

Os Osmium

76

192

Ir Iridium

77

195

Pt Platinum

78

197

Au Gold

79

201

Hg Mercury

80

204

Tl Thallium

81

207

Pb Lead

82

209

Bi Bismuth

83

Po Polonium

84

At Astatine

85

Rn Radon

86

Fr Francium

87

226

Ra Radium

88

227

Ac Actinium

89 +

*58-71 Lanthanoid series +90-103 Actinoid series

140

Ce Cerium

58

141

Pr Praseodymium

59

144

Nd Neodymium

60

Pm Promethium

61

150

Sm Samarium

62

152

Eu Europium

63

157

Gd Gadolinium

64

159

Tb Terbium

65

162

Dy Dysprosium

66

165

Ho Holmium

67

167

Er Erbium

68

169

Tm Thulium

69

173

Yb Ytterbium

70

175

Lu Lutetium

71

Keya

X b

a = relative atomic mass

X = atomic symbol b = proton (atomic) number

232

Th Thorium

90

Pa Protactinium

91

238

U Uranium

92

Np Neptunium

93

Pu Plutonium

94

Am Americium

95

Cm Curium

96

Bk Berkelium

97

Cf Californium

98

Es Einsteinium

99

Fm Fermium

100

Md Mendelevium

101

No Nobelium

102

Lr Lawrencium

103

The volume of one mole of any gas is 24 dm 3 at room temperature and pressure (r.t.p.)

Sec 2 Term 1_1 Notes

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