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Periodic Table Flashcards

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1
Q

Why were the elements arranged in the form of a table?

A

Because a need arose to group elements into families whose elements showed maximum resemblance.

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2
Q

What was observed when elements were arranged in the form of a table?

A

On an arrangement in the form of a table, properties of elements were seen to reappear at regular intervals in the table.

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3
Q

What basis of arrangement did early chemists use?

A

They arranged elements on basis of valency, metallic and non-metallic character

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4
Q

Why was the early chemists’ method discarded?

A

Method discarded as elements showed variable valency and dual character.

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5
Q

Name the scientist who arranged elements in increasing order of atomic weights and give the year.

A

Dobereiner in 1915 arranged elements in increasing order of their atomic weights.

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6
Q

What was Dobereiner’s method of arrangement?

A
  1. Elements were arranged into groups of 3 called triads.

2. Atomic weight of the middle element was found generally to be the average of the other two.

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7
Q

Why was Dobereiner’s method discarded?

A

Method discarded since it did not hold true for all elements.

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8
Q

Which scientist arranged elements in increasing order of atomic weights in groups of 8?

A

Newland, in 1864, arranged elements in increasing order of their atomic weights.

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9
Q

What was Newland’s method of arrangement?

A
  1. Elements were arranged in a series of 8.

2. Properties of every eight element was found to be a repetition of the 1st element.

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10
Q

Why was Newland’s method discarded?

A

Because it failed to leave spaces for undiscovered elements.

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11
Q

Name the scientist who arranged elements in increasing order of atomic weights into a table and give the year.

A

Mendeleeff, in 1869

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12
Q

What was Mendeleeff’s method of arrangement?

A
  1. Arranged elements in increasing order of their atomic weights.
  2. Elements were arranged in the form of a table called Mendeleeff’s Periodic Table.
  3. He stated that properties of elements were periodic functions of their atomic weights.
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13
Q

Why was Mendeleeff’s method discarded?

A

Because it could not justify the position of certain elements, rare earths and isotopes.

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14
Q

Which scientist arranged elements in increasing order of their atomic numbers?

A

Moseley in 1912

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15
Q

What was Moseley’s method of arrangement?

A
  1. Moseley arranged elements in increasing order of their atomic numbers.
  2. Elements were arranged in a modified table called modern periodic table.
  3. He stated that properties of elements were periodic functions of their atomic nos.
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16
Q

Why was Moseley’s method adopted?

A

This is because it removes most of the defects of Mendeleeff’s method i.e., it justifies the position of certain elements, rare earths and isotopes.

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17
Q

What reasons did Moseley give for modifying Mendeleeff’s periodic table?

A

Physical and and chemical properties of elements, depend on the no. of electrons and their arrangement & atomic no. is equal to the no. of electrons in the energy shells of an atom.

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18
Q

What did Moseley state which was opposite to Mendeleeff’s theory?

A

He stated that the basis of classification of elements be according to increasing atomic numbers and not atomic weights.

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19
Q

What is the fundamental property of an element?

A

Atomic no. is the fundamental property of an element.

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20
Q

Why is atomic no. the fundamental property of an element?

A

This is because the physical and chemical properties of an element depend on the no. of electrons and their arrangement, and atomic no. is equal to the no. of electrons in the energy shells of an atom.

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21
Q

What does the Modern Periodic Law state?

A

The MPL states that physical and chemical properties of elements are periodic functions of their at. no.

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22
Q

What is the arrangement of elements in the periodic table based on?

A

Based on the Modern Periodic Law.

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23
Q

Name the salient features of the Modern Periodic Table?

A
  1. Classification
  2. Position
  3. Methodical Arrangement
  4. Separation of Elements
  5. Periodicity of Elements
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24
Q

Elaborate on how classification is a salient feature of the MPT.

A

Period table is based on basic fundamental property - atomic no.

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25
Q

Elaborate on how position is a salient feature of the modern periodic table.

A

The MPT arranges elements in increasing the order of atomic numbers in seven horizontal rows called periods and eighteen vertical columns called groups.

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26
Q

Why is the completion of each period logical?

A

Since each period begins with an element having one electron in outermost shell and ends with zero group element having completely filled outermost shell.

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27
Q

What transition is seen across a period?

A

A transition from metallic to non-metallic character is seen across a period.

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28
Q

What is common between each vertical column?

A

Each vertical column accommodates elements with the same electronic configuration, thus having similar properties.

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29
Q

Which groups are called normal elements?

A

Groups 1, 2 & 13 to 17 [I Ato VII A] are called normal elements.

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30
Q

Which groups are called transition elements?

A

Groups 3 to 12 [I B to VII B & VIII] are called transition elements.

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31
Q

Which group contains noble or inert gases?

A

Group 18[zero] at extreme right contains noble gases.

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32
Q

Elaborate on the separation of metals in the MPT

A
  1. Reactive metals are placed in group 1[I A] and 2 [II A].
  2. Transition elements (metals) are placed in the middle.
  3. Non-metals are placed in the upper right corner of the periodic table.
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33
Q

What is periodicity of elements?

A

Gradual change in properties is seen with increases in atomic number in the periodic table.

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34
Q

What is periodicity in properties?

A

Recurrence in properties with elements belonging to same subgroup in the periodic table after a difference of 2,8, 18 or 32 in atomic numbers due to recurrence of similar valence shell electronic configuration is called periodicity in properties.

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35
Q

Define a period in the MPT.

A

A period is a horizontal row of elements arranged in increasing order of atomic numbers.

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36
Q

How many periods are there in the periodic table?

A

7 periods in the MPT

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37
Q

What does the period no. signify?

A

The period no. signifies the no. of electrons shells of an element.

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38
Q

Which are the short periods? Why?

A

Periods 1,2 and 3 are the short periods. Period 1 has only 2 elements, and periods 2 and 3 both have 8 elements.

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39
Q

Which are the long periods?

A

Periods 4,5,6 and 7 are the long periods with 18 , 18, 32 and 26 elements respectively.

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40
Q

In which period does the Lanthanide series fall?

A

In period 6

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41
Q

In which period does the actinide series fall?

A

In period 7

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42
Q

What are bridge elements?

A

Bridge elements are those which show similarities in properties diagonally with the period of the next group.

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43
Q

In which period are bridge elements found?

A

Bridge elements are found in period 2.

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44
Q

Name the bridge elements which show similarity with typical elements

A

Li, Be, B , C

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45
Q

Which typical element does Li sho similarity in properties with?

A

Li (atomic no. 3) shows similarity with Mg (at. no. 12)

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46
Q

Which typical element does Be show similarity in properties with?

A

Be (At. no. 4) shows similarity with Al (At. no. 13)

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47
Q

Which typical element does B show similarity in properties with?

A

B (at. no. 5) shows similarity with Si (at. no. 14).

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48
Q

What remains same across a period?

A

No. of electrons shells remains same across a period.

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49
Q

Which shells do electrons enter in period 1?

A

Electrons enter K shell in period 1

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50
Q

Which shell do electrons enter in period 2?

A

Electrons enter L shell in period 2.

M shell in period 3, and so on.

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51
Q

What increases across a period?

A

The no. of valence electrons increases by one across a period. Consequently, there is a transition from metallic to non-metallic character.

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52
Q

Non-metallic character _ across a period.

A

increases

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53
Q

Name the chlorides which are solid and have ionic bonding.

A
  1. Sodium chloride (NaCl)
  2. Magnesium chloride (MgCl2)
  3. Aluminium chloride (AlCl3)
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54
Q

Nam the chloride which is solid and goes through both ionic and covalent bonding.

A

Aluminium chloride (AlCl3) goes through both ionic and covalent bonding and is solid.

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55
Q

Name the chlorides which are liquid and go through covalent bonds.

A
  1. Silicon tetrachloride (SiCl4)
  2. Phosphorous trichloride
  3. Phosphorus pentachloride
  4. Disulphur dichloride
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56
Q

Name a chloride which can either be liquid or solid but goes through covalent bonding.

A

Phosphorus trichloride or phosphorus pentachloride

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58
Q

Name a basic oxide that goes through electrovalent bonding.

A

Magnesium oxide (MgO)

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59
Q

Name an amphoteric oxide that goes through electrovalent bonding.

A

Aluminium oxide (Al2O3)

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60
Q

Name a weakly acidic oxide that goes through covalent bonding.

A

silicon dioxide (SiO2)

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61
Q

Name three acidic oxides that go through covalent bonding.

A 
Phosphorous pentoxide (P2O5)
Sulphur dioxide (SO2)
Sulphur trioxide (SO3)
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62
Q

Name a strongly acidic oxide that goes through covalent bonding.

A

Chlorine heptoxide (Cl2O7)

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63
Q

Name a hydroxide that is a strong base.

A

Sodium hydroxide (NaOH)

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64
Q

Name a hydroxide that is a weak base.

A

Magnesium hydroxide (Mg(OH)2)

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65
Q

Name a hydroxide that is amphoteric

A

Aluminium hydroxide (Al(OH3))

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66
Q

What is an oxyacid?

A

An oxyacid is an acid that contains oxygen. Specifically, it is a compound that contains hydrogen, oxygen, and at least one other element.

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67
Q

Name three oxyacids which is a weak acid.

A
  1. Metasilicic acid (H₂SiO₃)
  2. Metaphosphoric Acid (HPO3
  3. Orthophosphoric acid ( H₃PO ₄)
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68
Q

Name three oxyacids that are strong acids.

A
  1. Sulphurous acid (H₂SO₃)
  2. Sulphuric acid (H₂SO₄)
  3. Perchloric acid (HClO₄)
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69
Q

What is a hydride?

A

In compounds that are regarded as hydrides, the hydrogen atom is bonded to a more electropositive element or group.

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70
Q

Name a hydride that is a strong base.

A

Sodium hydride (NaH)

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71
Q

Name a hydride that is a weak base.

A

Magnesium hydride (MgH₂)

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72
Q

Name three hydrides that are weaker bases.

A
  1. Aluminium hydride (AlH₃)
  2. Silane (SiH₄)
  3. Phosphine (PH₃)
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73
Q

Name a hydride that is a weak acid.

A

Hydrogen sulphide (H ₂S)

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74
Q

Name a hydride that is a strong acid.

A

Hydrochloric acid (HCl)

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75
Q

Name a strongly basic oxide that goes through electrovalent bonding.

A

Sodium oxide (Na2O)

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75
Q

What transition is seen across period 2 with reference to the valency?

A

The valency of the elements in period 2 increases from 1 to 4, 1 being a metal Lithium, 3 being a metalloid Boron, 4 being a non-metal Carbon. Then it decreases from 3 to 0, 3 being nitrogen and 0 being neon.

Increases from 1 to 4, then 3 to 0

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76
Q

Name the metals present in period 2.

A

Lithium and Beryllium

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77
Q

Name the non-metals present in period 2.

A

Carbon, Nitrogen, Oxygen, Fluorine

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78
Q

Name the metalloid present in period 2

A

Boron

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79
Q

Name the noble gas present in period 2.

A

Neon

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80
Q

Name the two elements present in period 1.

A

Hydrogen [H-group 1(I A)] & Helium [ 2He - group 18 (0)]

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81
Q

Name the metals present in period 3 with their valency.

A

Sodium - Valency 1
Magnesium - Valency 2
Aluminium - Valency 3

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82
Q

Name the metalloid in period 3, along with valency.

A

Silicon - Valency 4

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83
Q

Name the non-metals in period 3.

A

Phosphorus - Valency 3
Sulphur - Valency 2
Chlorine - Valency 1

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84
Q

Which is the noble gas in period 3?

A

Argon - Valency 0

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85
Q

How many groups in the periodic table?

A

There are eighteen vertical columns - [with eight main groups] in the periodic table.

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86
Q

What does the group no. signify?

A

Group no. signifies the no. of valence electrons of an element, which are electrons present in the outermost shell of an element.

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87
Q

How many valence electrons do transition elements have?

A

2 valence electrons

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88
Q

Which type of elements does group 1 contain?

A

Group 1 [I A] contains alkali metals.

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89
Q

Which elements does group 1[I A] contain?

A

(Light metals) Li (At. no. 3) to Francium (At. no. 87)

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90
Q

Which type of elements does group 2 [II A] contain?

A

Group 2[II A] contains alkaline earth metals.

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91
Q

Which elements does group 2[II A] contain?

A

(Light metals) From Beryllium (At. No. 4) to Radium (At. No. 88)

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92
Q

What type of elements do group 3 to 12 contain?

A

Groups 3 to 12 contain transition elements.

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93
Q

Which are the transition elements & which groups are they in?

A

The transition elements are Scandium (At. No. 21) to Zinc (At. No. 30) in groups 3 to 12. These are heavy metals.

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94
Q

Which are the inner transition elements & which groups are they in?

A

The inner transition elements range from Yttrium [Y] (At No. 39) to Cadmium [Cd] (At. No. 48). They are in the groups IB to VIIB, VIII

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95
Q

Which series is found in period 6 and groups 3 to 12?

A

The lanthanide series is found in period 6. It ranges from the elements Cerium [Ce] (At. No. 58) to Lutetium [Lu] (At. No. 71)

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96
Q

What is the sequence of elements in period 6 and groups 3 to 12?

A

Lanthanum [La] (At. No. 57); Ce (58) to Lu (71); Hafnium [Hf] (At. No. 72) to Mercury [Hg] (80).

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97
Q

Which series is found in period 7 and groups 3 to 12?

A

The actinide series is found in period 7. It ranges from the elements Thorium [Th] (At. No. 90) to Lawrencium [Lr] (At. No. 103)

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98
Q

What is the sequence of elements in period 7 and groups 3 to 12?

A

Actinium [Ac] (At. No. 89); Th (90) to Le (103); Rutherfordium[Rf] (At. No. 104) to Copernicium [Cn] (112).

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99
Q

What type of elements do group 13 to 16 contain?

A

They contain post transition elements, from IIIA to VIA.

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100
Q

What is the range of elements in groups 13 to 16?

A

Aluminium [Al] (13) to Thallium [Ti] (81)
Germanium [Ge] (32) to Lead [Pb] (82)
Antimony [Sb] (51) to Bismuth [Bi] (83)
Polonium [Po] (84)

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101
Q

What type of elements are found in group 17 [VII A]?

A

Halogens

102
Q

What is the sequence of elements in group 17?

A

Fluorine [F] (9) to Astatine [At] (85)

103
Q

What type of elements are found in group 18 [0]?

A

Noble/inert gases

104
Q

What is the sequence of elements in group 18?

A

Helium [He] (2) to Radon [Rn] (86}

105
Q

Where do transition elements lie in the periodic table?

A

They lie between strongly electropositive metals on the left & the least electropositive elements on the right.

106
Q

Why do transition elements have similar properties?

A

This is because they have the same valence electrons and electron change occurs in the inner orbitals.

107
Q

Where do the inner transition elements lie?

A

These are two horizontal rows at the bottom of the table.

108
Q

Which two series do inner transition elements form?

A

They form the lanthanide [rare earths] and the actinide [radio active] series of 14 elements each.

109
Q

Which properties remain same going down a subgroup?

A
  1. Valence electrons of elements in a sub-group remain same.

2. Chemical properties of elements in a sub-group remain the same.

110
Q

Why do chemical properties of elements in a sub-group remain the same?

A

Because chemical properties are dependent on outer electronic configuration?

111
Q

Which properties change as you go down a subgroup?

A
  1. Metallic character (electropositive char.) increases.

2. No. of electron shells increases by one.

112
Q

Name the elements in group 1 [I A]

A

Lithium, sodium, potassium, rubedium, caesium, francium

113
Q

What is the valency of elements in group 1 [I A]?

A

Univalent [1 valence electron]

114
Q

Give 3 points about the nature of elements in group 1

A
  1. Highly reactive
  2. Highly electropositive
  3. Light, soft metals
    (metallic in nature)
115
Q

Why can group 1 elements be cut with a knife?

A

This is because they have only one single electron in their valence shell and the metallic bonding is weak between the atoms. We can say that these metals have weak binding energy in the crystal lattice, hence they are easy to cut.

116
Q

What kind of conductivity do elements of group 1 have?

A

They are good conductors of heat and electricity.

117
Q

What is the nature of group 1 elements with reference to reduction/oxidization?

A

They are strong reducing agents.

118
Q

Why are alkali metals called electron donors?

A

Alkali metals

119
Q

Why are alkali metals strong reducing agents?

A

This is because they are electron donors. They have one valence electron which is easily removed from the outer shell.

120
Q

What is the nature of group 1 elements with reference to electronegativity?

A

They have low electronegativity (electropositive character increases from Li to Cs)

121
Q

What is formed when group 1 elements react with non-metals?

A

Electrovalent compounds formed

122
Q

Give examples of electrovalent compounds formed when group 1 elements react with non-metals.

A

Sodium chloride (NaCl), potassium bromide (KBr)

123
Q

What is formed when group 1 elements react with hydrogen?

A

Ionic hydrides formed

124
Q

Give examples of ionic hydrides formed when alkali metals react with hydrogen.

A

Lithium Hydride (LiH), Sodium hydride (NaH)

125
Q

Name the elements in group 17 [VII A]

A

Fluorine, chlorine, bromine, iodine, astatine

126
Q

What is the valency of elements in group 17 [VII A]?

A

Univalent [7 valence electrons]

127
Q

Give 3 points about the nature of elements in group 17

A
  1. Highly reactive
  2. Highly electronegative
  3. Non-metals
    (non-metallic in nature)
128
Q

Which elements in group 17 are gaseous?

A

Fluorine and Chlorine

129
Q

Which elements in group 17 are liquid?

A

Bromine

130
Q

Which elements in group 17 are solid at room temp?

A

Iodine

131
Q

What kind of conductivity do elements of group 17 have?

A

Bad or non-conductors of heat and electricity

132
Q

What is the nature of group 17 elements with reference to reduction/oxidization?

A

They are strong oxidizing agents.

133
Q

Why are the elements of group 17 strong oxidizing agents?

A

This is because they are electron acceptors as they have 7 valence electrons and need only 1 to attain stable electronic configuration.

134
Q

What is the nature of group 17 elements with reference to electronegativity?

A

They have high electronegativity. Electronegative character decreases from Fluorine to Iodine.

135
Q

What is formed when group 17 elements react with non-metals?

A

Covalent compounds are formed.

136
Q

Give examples of covalent compounds formed when group 17 elements react with non-metals.

A

Hydrochloric acid (HCl), Phosphorus trichloride (PCl₃), Disulfur dichloride (S₂Cl₂)

137
Q

What is formed when group 17 elements react with hydrogen?

A

Covalent hydrides are formed when group 17 elements react with hydrogen.

138
Q

Give examples of covalent hydrides formed when group 17 elements react with hydrogen.

A

Hydrogen Fluoride (HF), Hydrochloric acid (HCl), Hydrogen Bromide (HBr), Hydrogen Iodide (HI)

139
Q

What does periodicity in properties of elements mean?

A

Occurrence of characteristic properties of elements at definite intervals in the modern periodic table when elements are arranged in increasing order of their atomic numbers.

140
Q

When do definite intervals occur in the MPT?

A

The definite intervals are after difference of either 2 or 8 or 18 or 32 in atomic numbers.

141
Q

What are periodic properties?

A

The properties which appear at regular intervals in the periodic table are called periodic properties.

142
Q

Name the periodic properties found in the MPT.

A
  1. Atomic radii
  2. Ionisation potential
  3. Electron affinity
  4. Electronegativity
  5. Non-metallic and metallic character
  6. Density
  7. Melting and Boiling point
  8. Nature of oxides, hydrides, and oxyacids.
143
Q

What is the reason for periodicity in properties in periods & groups?

A

After definite intervals of atomic no., similar valence shell electronic configuration occurs. Properties of elements depend on the no. & arrangement of electrons in various shells including valence shells. In the same period or subgroup, increase or decrease in a particular property is due to the gradual change in electronic configuration in the arranged elements.

144
Q

What does atomic radius mean?

A

It is the distance between the centre of the nucleus and the outermost shell of the atom.

145
Q

What does ionisation potential mean?

A

It is the amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom.

146
Q

What does electron affinity mean?

A

It is the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.

147
Q

What does electronegativity mean?

A

It is the tendency of an atom to attract electrons to itself when combined in a compound

148
Q

What does metallic character mean?

A

In terms of electron loss or gain, an element is a metal if it gains one or more electrons

149
Q

What does non-metallic character mean?

A

In terms of electron loss or gain, an element is a non-metal if it loses one or more electrons

150
Q

What is the unit of atomic radii?

A

Angstron unit = Å

151
Q

What is bond length?

A

Bond length or bond distance is defined as the average distance between nuclei of two bonded atoms in a molecule.

152
Q

Which factors affect the atomic size?

A
  1. Number of shells

2. Nuclear charge

153
Q

What effect does no. of shells have on atomic size?

A

Atomic size increases if no. of shells increases

154
Q

Why does atomic size increase on increase of no. of shells?

A

This is because when no. of shells increases, the distance of the outermost shell from the nucleus increases.

155
Q

What effect does nuclear charge have on atomic size?

A

Atomic size decreases when nuclear charge increases.

156
Q

Why does atomic size decrease with increase in nuclear charge?

A

When nuclear charge increases, the electrons in the outermost shell are attracted with increasing force.

157
Q

What is nuclear charge?

A

It is the positive charge on the nucleus of an atom.

158
Q

What is nuclear charge equivalent to?

A

It is equivalent to the atomic no. of an atom.

159
Q

What is the trend in atomic size across a period?

A

Atomic size decreases across a period left to right.

160
Q

What remains the same across a period?

A

The no. of shells remains the same, and therefore the atomic size remains unaffected

161
Q

What increases across a period?

A

Nuclear charge increases across a period, and therefore, the atomic size decreases.

162
Q

Which elements have the largest and smallest atomic radius in period 2?

A

Lithium has the largest atomic radius and Fluorine has the smallest atomic radius.

163
Q

Why does Neon have a larger atomic radius despite being the last in the period?

A

In inert gases, the outer shell is completely filled resulting in a force of repulsion. The effect of the nuclear pull over the valence shell electrons is not seen.

164
Q

How is a cation formed?

A

It is formed by loss of electron. Na - 1e- –> Na 1+

165
Q

Why is a cation smaller than the parent neutral atom?

A

The remaining electrons in the cation are strongly attracted by the nucleus, thus decreasing the cation size.

166
Q

What is the trend in atomic size down a group?

A

Atomic size (radii) increases down a group

167
Q

What increases down a group?

A

No. of shells increases down a group, therefore atomic size increases.

168
Q

Why does no. of shells increase down a group?

A

New shells are added with increasing atomic no.

169
Q

Why does atomic size increase on increase in no. of shells down a group?

A

Nuclear distance from valence electrons increases on increase in no. of shells down a group

170
Q

What happens to nuclear charge down a group?

A

Nuclear charge increases down a group

171
Q

Why does nuclear charge increase down a group?

A

Nuclear charge is equivalent to the atomic no. of an element, and it increases as atomic no. increases down a group

172
Q

What happens to at. size as nuclear charge increases down a group?

A

Atomic size should decrease as nuclear charge increases down a group.

173
Q

Why does overall atomic radius increases down a subgroup?

A

Increase in no. of shells dominates over increase in nuclear charge, therefore overall atomic radius increases.

174
Q

What is the first ionisation potential?

A

Energy required to remove 1st electron is called first I.P.

M (atom) -> M+ (ion) + e- (electron)

175
Q

What is the second ionisation potential?

A

Energy required to remove second electron is called second I.P.
M+ -> M++ + e-

176
Q

Which I.P. is more?

A

The second I.P. is more.

177
Q

What is the unit of I.P.?

A

Electron volt = e V

178
Q

Name the factors which affect the I.P.

A
  1. Atomic Size

2. Nuclear charge

179
Q

How does atomic size affect ionisation potential?

A

When atomic size increases, I.P. decreases.

180
Q

Why does I.P. decrease when atomic size increases?

A

When atomic size increases, the nuclear attraction on the outer electrons decreases. hence, the outer electrons are loosely held. Therefore, the I.P. decreases.

181
Q

How does nuclear charge affect I.P.?

A

When nuclear charge increases, I.P. increases.

182
Q

Why does I.P. increase when nuclear charge increases?

A

When nuclear charge increases. the nuclear attraction on the outer electrons increases. Hence, the puter electrons are more firmly held. Therefore. I.P. increases.

183
Q

What is the trend in I.P. across a period from left to right?

A

I.P. increases across a period from left to right.

184
Q

How do the factors affecting I.P. vary across a period?

A

The atomic radii decreases, so I.P. increases.

The nuclear charge increases, so I.P. increases.

185
Q

Which element has the highest I.P.?

A

The element helium [He] has the highest I.P.

186
Q

Which element has the lowest I.P.?

A

The element caesium [Cs] has the lowest I.P.

187
Q

Why do metals have low I.P. as compared to non-metals?

A

Metals lose electrons and thus, have a lower I.P. than non-metals.

188
Q

Why does francium not have the lowest I.P. despite being in the last period?

A

Francium [Fr] is radioactive.

189
Q

What is the trend in I.P. down a group?

A

I.P. decreases down a group

190
Q

Why does overall I.P. decrease down a group?

A

Increase in atomic radii dominates over increase in nuclear charge. Therefore, overall I.P. decreases down a group.

191
Q

How do the factors affecting I.P. vary down a group?

A

Atomic radii increases down a group -> I.P. decreases

Nuclear charge increases -> I.P. should increase

192
Q

What is the 1st electron affinity?

A

The energy liberated when an atom in the gaseous state accepts an electron to form an anion.

193
Q

What is the unit of electron affinity?

A

Electron volt = e V

194
Q

With which sign is electron affinity represented?

A

Electron affinity is represented by a negative sign

195
Q

Name the factors which affect electron affinity

A
  1. Atomic size

2. Nuclear charge

196
Q

How does atomic size affect electron affinity?

A

When atomic size increases, electron affinity decreases.

197
Q

Why does electron affinity decrease on increase of atomic size?

A

Electron affinity is the tendency of an atom to accept electrons. A small atom takes up electrons more readily than a large atom since nucleus has greater attraction on the electrons.

198
Q

How does nuclear charge affect electron affinity?

A

When nuclear charge increases, electron affinity increases.

199
Q

Why does e.a. increase on increase of nuclear charge?

A

Increase in nuclear charge increases the tendency of an atom to accept electrons.

200
Q

What is the trend in electron affinity across a period with relation to a.r and n.c?

A

Atomic radii decreases - Electron affinity increases

Nuclear charge increases - Electron affinity increases

201
Q

What is the e.a highest and lowest for?

A

Highest for halogens [group 17]

Least for alkali metals [group 1]

202
Q

Why does neon have e.a zero?

A

Inert gases with stable electronic configuration find it difficult to accept electrons.

203
Q

Why do inert gases not form ions?

A

They do not form ions since their outermost shell is completely filled. They need not accept or donate any electrons since they are already stable. They have no urge to destabilize themselves through gain or loss of electrons and hence do not form ions.

204
Q

What is the trend in e.a across a period?

A

E.A. increases across a period from left to right

205
Q

What is the trend in electron affinity down a group with relation to a.r and n.c?

A

Atomic radii increases - Electron affinity decreases

Nuclear charge increases - Electron affinity should increase

206
Q

Why does e.a. decrease down a group?

A

Increase in atomic radii dominates over increase in nuclear charge. Hence, overall e.a. decreases

207
Q

What is the trend in electron affinity down a group?

A

E.A. decreases down a group

208
Q

What is the effect of E.A. on the electronegative/oxidizing nature of an element?

A

Greater the value of E.A., more electronegative, or more oxidizing is the element

209
Q

Which element has greater oxidization potential - electropositive or electronegative?

A

More electronegative element

210
Q

Define electronegativity

A

The tendency of an atom to attract electrons to itself when combined in a compound

211
Q

How does electronegativity affect the nature of bond?

A

Ionic bond is formed between combining atoms if atoms differ widely in electronegativity, whereas
Covalent bond is formed between combining atoms if atoms have nearly similar electronegativity

212
Q

Name the factors which affect E.N.?

A
  1. Atomic size

2. Nuclear charge

213
Q

How does atomic size affect E.N.?

A

When atomic size increases, E.N. decreases.

214
Q

Why does E.N. decrease on increase of atomic size?

A

Down a group, the atomic size increases, hence the number of energy levels (n) increases, and so does the distance between the nucleus and the outermost orbital. The increased distance and the increased shielding weaken the nuclear attraction, and so an atom can’t attract electrons as strongly.

Even though the nuclear charge increases when going down as well, the increase in radius is more important.

215
Q

How does nuclear charge affect E.N.?

A

Nuclear charge increases - E.N. increases

216
Q

Why does E.N. increase on increase of nuclear charge?

A

When nuclear charge increases, the tendency of an atom to accept electrons increases

217
Q

Name the factors affecting E.N.

A
  1. Atomic size

2. Nuclear charge

218
Q

How do the factors affecting E.N. vary across a period?

A

Atomic radii decreases -> E.N. increases

Nuclear charge increases -> E.N. increases

219
Q

Which elements have high electronegativity?

A

Non-metallic elements

220
Q

What is the trend in non-metallic character across a period?

A

It increases from left to right across a period

221
Q

Which is the most E.N. element in the periodic table?

A

Fluorine

222
Q

Which is the least E.N. element in the periodic table?

A

Caesium

223
Q

Who do noble gases have zero E.N.?

A

They have a complete octet and hence do not attract electrons to themselves

224
Q

What is the trend in overall E.N. across a period?

A

It increases across a period - left to right

225
Q

Which elements is E.N. highest for?

A

Halogens

226
Q

How do the factors affecting E.N. vary down a group?

A

Atomic radii increases -> E.N. decreases

Nuclear charge increases -> E.N. should increase

227
Q

Why does overall E.N. decrease down a group?

A

Increase in atomic radii dominates over increase in nuclear charge. Therefore, overall E.N. decreases

228
Q

Define metallic character

A

An atom is said to be a metal if it loses one or more electrons when supplied with energy

229
Q

Define non-metallic character

A

An atom is said to be a non-metal if it gains one or more electrons when supplied with energy

230
Q

Name the factors influencing M.C/N.M.C

A
  1. Atomic radii

2. Ionisation potential

231
Q

How does atomic radii influence M.C/N.M.C?

A

A.R. increases - Metallic character increases

A.R. increases - Non-metallic character decreases

232
Q

How does I.P. influence M.C/N.M.C?

A

I.P. increases - Metallic character decreases

I.P. increases - Non-metallic character increases

233
Q

Give characteristics of metallic atoms

A

They are present on the left side of the periodic table, have large atomic radii and high I.P. value, and tend to lose electrons

234
Q

Give characteristics of non-metallic atoms

A

They are present on the right side of the periodic table, have small atomic radii and high I.P. value, and tend to gain electrons

235
Q

How does the reactivity of an element depend on its tendency to lose or gain electrons?

A

Greater the tendency to lose electrons, the greater is the reactivity of the metal. Greater the tendency to gain electrons, the greater is the reactivity of the non-metal

236
Q

Are metals good reducing or oxidizing agents?

A

Metals are good reducing agents

237
Q

Are non-metals good reducing or oxidizing agents?

A

Non-metals are good oxidizing agents

238
Q

Give the trend in character across a period

A

Metallic character decreases across a period

Non-metallic character increases across a period

239
Q

How do the factors affecting character vary across a period?

A

A.R. decreases -> MC decreases, NMC increases

I.P. increases -> MC decreases, NMC increases

240
Q

How do the factors affecting character vary down a group?

A

A.R. increases -> MC increases, NMC decreases

I.P. decreases -> MC increases, NMC decreases

241
Q

Give the characteristics of the elements at the bottom of a group

A

Elements at the bottom of a group are most metallic, have large atomic size, lowest I.P., electrons are thus loosely held and will form ions from metals most readily and thus are most reactive

242
Q

Why do elements at the bottom of a group form ions from metals most readily/ are most reactive?

A

Thye have a large atomic size, the lowest I.P. and thus, electrons are loosely held

243
Q

Give the trend in character down a group

A

Metallic character increases down a group

Non-metallic character decreases down a group

244
Q

Give the formula for a metal atom losing an electron

A

M -> M+ (ion) + e- (electron)

245
Q

Give the formula for a non-metal atom gaining an electron

A

N + e- -> N- (ion)

246
Q

What are light metals?

A

Elements arranged in the periodic table having a neutron/proton ratio around 1 are stable elements

247
Q

Which light metals are stable elements?

A

light metals Na and K

248
Q

What are heavy metals?

A

Elements with an n/p ratio above 1.5 are considered radioactive unstable elements

249
Q

Give an example of a heavy metal that is a radioactive unstable element

A

Uranium

250
Q

Give the steps to calculating if an element is stable or unstable

A
  1. No. of protons (p) = Atomic no. (Z)
  2. No. of neutrons (n) = Mass no. (A) - Atomic no. (Z)
  3. n/p < 1.5 = light/stable, n/p > 1.5 = heavy/unstable

Chemistry (4 decks)

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