Periodic Table Flashcards
(250 cards)
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.
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.
3
Q
What basis of arrangement did early chemists use?
A
They arranged elements on basis of valency, metallic and non-metallic character
4
Q
Why was the early chemists’ method discarded?
A
Method discarded as elements showed variable valency and dual character.
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.
6
Q
What was Dobereiner’s method of arrangement?
A
- 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.
7
Q
Why was Dobereiner’s method discarded?
A
Method discarded since it did not hold true for all elements.
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.
9
Q
What was Newland’s method of arrangement?
A
- Elements were arranged in a series of 8.
2. Properties of every eight element was found to be a repetition of the 1st element.
10
Q
Why was Newland’s method discarded?
A
Because it failed to leave spaces for undiscovered elements.
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
12
Q
What was Mendeleeff’s method of arrangement?
A
- Arranged elements in increasing order of their atomic weights.
- Elements were arranged in the form of a table called Mendeleeff’s Periodic Table.
- He stated that properties of elements were periodic functions of their atomic weights.
13
Q
Why was Mendeleeff’s method discarded?
A
Because it could not justify the position of certain elements, rare earths and isotopes.
14
Q
Which scientist arranged elements in increasing order of their atomic numbers?
A
Moseley in 1912
15
Q
What was Moseley’s method of arrangement?
A
- Moseley arranged elements in increasing order of their atomic numbers.
- Elements were arranged in a modified table called modern periodic table.
- He stated that properties of elements were periodic functions of their atomic nos.
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.
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.
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.
19
Q
What is the fundamental property of an element?
A
Atomic no. is the fundamental property of an element.
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.
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.
22
Q
What is the arrangement of elements in the periodic table based on?
A
Based on the Modern Periodic Law.
23
Q
Name the salient features of the Modern Periodic Table?
A
- Classification
- Position
- Methodical Arrangement
- Separation of Elements
- Periodicity of Elements
24
Q
Elaborate on how classification is a salient feature of the MPT.
A
Period table is based on basic fundamental property - atomic no.
25
Elaborate on how position is a salient feature of the modern periodic table.
The MPT arranges elements in increasing the order of atomic numbers in seven horizontal rows called periods and eighteen vertical columns called groups.
26
Why is the completion of each period logical?
Since each period begins with an element having one electron in outermost shell and ends with zero group element having completely filled outermost shell.
27
What transition is seen across a period?
A transition from metallic to non-metallic character is seen across a period.
28
What is common between each vertical column?
Each vertical column accommodates elements with the same electronic configuration, thus having similar properties.
29
Which groups are called normal elements?
Groups 1, 2 & 13 to 17 [I Ato VII A] are called normal elements.
30
Which groups are called transition elements?
Groups 3 to 12 [I B to VII B & VIII] are called transition elements.
31
Which group contains noble or inert gases?
Group 18[zero] at extreme right contains noble gases.
32
Elaborate on the separation of metals in the MPT
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.
33
What is periodicity of elements?
Gradual change in properties is seen with increases in atomic number in the periodic table.
34
What is periodicity in properties?
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.
35
Define a period in the MPT.
A period is a horizontal row of elements arranged in increasing order of atomic numbers.
36
How many periods are there in the periodic table?
7 periods in the MPT
37
What does the period no. signify?
The period no. signifies the no. of electrons shells of an element.
38
Which are the short periods? Why?
Periods 1,2 and 3 are the short periods. Period 1 has only 2 elements, and periods 2 and 3 both have 8 elements.
39
Which are the long periods?
Periods 4,5,6 and 7 are the long periods with 18 , 18, 32 and 26 elements respectively.
40
In which period does the Lanthanide series fall?
In period 6
41
In which period does the actinide series fall?
In period 7
42
What are bridge elements?
Bridge elements are those which show similarities in properties diagonally with the period of the next group.
43
In which period are bridge elements found?
Bridge elements are found in period 2.
44
Name the bridge elements which show similarity with typical elements
Li, Be, B , C
45
Which typical element does Li sho similarity in properties with?
Li (atomic no. 3) shows similarity with Mg (at. no. 12)
46
Which typical element does Be show similarity in properties with?
Be (At. no. 4) shows similarity with Al (At. no. 13)
47
Which typical element does B show similarity in properties with?
B (at. no. 5) shows similarity with Si (at. no. 14).
48
What remains same across a period?
No. of electrons shells remains same across a period.
49
Which shells do electrons enter in period 1?
Electrons enter K shell in period 1
50
Which shell do electrons enter in period 2?
Electrons enter L shell in period 2.
| M shell in period 3, and so on.
51
What increases across a period?
The no. of valence electrons increases by one across a period. Consequently, there is a transition from metallic to non-metallic character.
52
Non-metallic character _ across a period.
increases
53
Name the chlorides which are solid and have ionic bonding.
1. Sodium chloride (NaCl)
2. Magnesium chloride (MgCl2)
3. Aluminium chloride (AlCl3)
54
Nam the chloride which is solid and goes through both ionic and covalent bonding.
Aluminium chloride (AlCl3) goes through both ionic and covalent bonding and is solid.
55
Name the chlorides which are liquid and go through covalent bonds.
1. Silicon tetrachloride (SiCl4)
2. Phosphorous trichloride
3. Phosphorus pentachloride
4. Disulphur dichloride
56
Name a chloride which can either be liquid or solid but goes through covalent bonding.
Phosphorus trichloride or phosphorus pentachloride
58
Name a basic oxide that goes through electrovalent bonding.
Magnesium oxide (MgO)
59
Name an amphoteric oxide that goes through electrovalent bonding.
Aluminium oxide (Al2O3)
60
Name a weakly acidic oxide that goes through covalent bonding.
silicon dioxide (SiO2)
61
Name three acidic oxides that go through covalent bonding.
```
Phosphorous pentoxide (P2O5)
Sulphur dioxide (SO2)
Sulphur trioxide (SO3)
```
62
Name a strongly acidic oxide that goes through covalent bonding.
Chlorine heptoxide (Cl2O7)
63
Name a hydroxide that is a strong base.
Sodium hydroxide (NaOH)
64
Name a hydroxide that is a weak base.
Magnesium hydroxide (Mg(OH)2)
65
Name a hydroxide that is amphoteric
Aluminium hydroxide (Al(OH3))
66
What is an oxyacid?
An oxyacid is an acid that contains oxygen. Specifically, it is a compound that contains hydrogen, oxygen, and at least one other element.
67
Name three oxyacids which is a weak acid.
1. Metasilicic acid (H₂SiO₃)
2. Metaphosphoric Acid (HPO3
3. Orthophosphoric acid ( H₃PO ₄)
68
Name three oxyacids that are strong acids.
1. Sulphurous acid (H₂SO₃)
2. Sulphuric acid (H₂SO₄)
3. Perchloric acid (HClO₄)
69
What is a hydride?
In compounds that are regarded as hydrides, the hydrogen atom is bonded to a more electropositive element or group.
70
Name a hydride that is a strong base.
Sodium hydride (NaH)
71
Name a hydride that is a weak base.
Magnesium hydride (MgH₂)
72
Name three hydrides that are weaker bases.
1. Aluminium hydride (AlH₃)
2. Silane (SiH₄)
3. Phosphine (PH₃)
73
Name a hydride that is a weak acid.
Hydrogen sulphide (H ₂S)
74
Name a hydride that is a strong acid.
Hydrochloric acid (HCl)
75
Name a strongly basic oxide that goes through electrovalent bonding.
Sodium oxide (Na2O)
75
What transition is seen across period 2 with reference to the valency?
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
76
Name the metals present in period 2.
Lithium and Beryllium
77
Name the non-metals present in period 2.
Carbon, Nitrogen, Oxygen, Fluorine
78
Name the metalloid present in period 2
Boron
79
Name the noble gas present in period 2.
Neon
80
Name the two elements present in period 1.
Hydrogen [H-group 1(I A)] & Helium [ 2He - group 18 (0)]
81
Name the metals present in period 3 with their valency.
Sodium - Valency 1
Magnesium - Valency 2
Aluminium - Valency 3
82
Name the metalloid in period 3, along with valency.
Silicon - Valency 4
83
Name the non-metals in period 3.
Phosphorus - Valency 3
Sulphur - Valency 2
Chlorine - Valency 1
84
Which is the noble gas in period 3?
Argon - Valency 0
85
How many groups in the periodic table?
There are eighteen vertical columns - [with eight main groups] in the periodic table.
86
What does the group no. signify?
Group no. signifies the no. of valence electrons of an element, which are electrons present in the outermost shell of an element.
87
How many valence electrons do transition elements have?
2 valence electrons
88
Which type of elements does group 1 contain?
Group 1 [I A] contains alkali metals.
89
Which elements does group 1[I A] contain?
(Light metals) Li (At. no. 3) to Francium (At. no. 87)
90
Which type of elements does group 2 [II A] contain?
Group 2[II A] contains alkaline earth metals.
91
Which elements does group 2[II A] contain?
(Light metals) From Beryllium (At. No. 4) to Radium (At. No. 88)
92
What type of elements do group 3 to 12 contain?
Groups 3 to 12 contain transition elements.
93
Which are the transition elements & which groups are they in?
The transition elements are Scandium (At. No. 21) to Zinc (At. No. 30) in groups 3 to 12. These are heavy metals.
94
Which are the inner transition elements & which groups are they in?
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
95
Which series is found in period 6 and groups 3 to 12?
The lanthanide series is found in period 6. It ranges from the elements Cerium [Ce] (At. No. 58) to Lutetium [Lu] (At. No. 71)
96
What is the sequence of elements in period 6 and groups 3 to 12?
Lanthanum [La] (At. No. 57); Ce (58) to Lu (71); Hafnium [Hf] (At. No. 72) to Mercury [Hg] (80).
97
Which series is found in period 7 and groups 3 to 12?
The actinide series is found in period 7. It ranges from the elements Thorium [Th] (At. No. 90) to Lawrencium [Lr] (At. No. 103)
98
What is the sequence of elements in period 7 and groups 3 to 12?
Actinium [Ac] (At. No. 89); Th (90) to Le (103); Rutherfordium[Rf] (At. No. 104) to Copernicium [Cn] (112).
99
What type of elements do group 13 to 16 contain?
They contain post transition elements, from IIIA to VIA.
100
What is the range of elements in groups 13 to 16?
Aluminium [Al] (13) to Thallium [Ti] (81)
Germanium [Ge] (32) to Lead [Pb] (82)
Antimony [Sb] (51) to Bismuth [Bi] (83)
Polonium [Po] (84)
101
What type of elements are found in group 17 [VII A]?
Halogens
102
What is the sequence of elements in group 17?
Fluorine [F] (9) to Astatine [At] (85)
103
What type of elements are found in group 18 [0]?
Noble/inert gases
104
What is the sequence of elements in group 18?
Helium [He] (2) to Radon [Rn] (86}
105
Where do transition elements lie in the periodic table?
They lie between strongly electropositive metals on the left & the least electropositive elements on the right.
106
Why do transition elements have similar properties?
This is because they have the same valence electrons and electron change occurs in the inner orbitals.
107
Where do the inner transition elements lie?
These are two horizontal rows at the bottom of the table.
108
Which two series do inner transition elements form?
They form the lanthanide [rare earths] and the actinide [radio active] series of 14 elements each.
109
Which properties remain same going down a subgroup?
1. Valence electrons of elements in a sub-group remain same.
| 2. Chemical properties of elements in a sub-group remain the same.
110
Why do chemical properties of elements in a sub-group remain the same?
Because chemical properties are dependent on outer electronic configuration?
111
Which properties change as you go down a subgroup?
1. Metallic character (electropositive char.) increases.
| 2. No. of electron shells increases by one.
112
Name the elements in group 1 [I A]
Lithium, sodium, potassium, rubedium, caesium, francium
113
What is the valency of elements in group 1 [I A]?
Univalent [1 valence electron]
114
Give 3 points about the nature of elements in group 1
1. Highly reactive
2. Highly electropositive
3. Light, soft metals
(metallic in nature)
115
Why can group 1 elements be cut with a knife?
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
What kind of conductivity do elements of group 1 have?
They are good conductors of heat and electricity.
117
What is the nature of group 1 elements with reference to reduction/oxidization?
They are strong reducing agents.
118
Why are alkali metals called electron donors?
Alkali metals
119
Why are alkali metals strong reducing agents?
This is because they are electron donors. They have one valence electron which is easily removed from the outer shell.
120
What is the nature of group 1 elements with reference to electronegativity?
They have low electronegativity (electropositive character increases from Li to Cs)
121
What is formed when group 1 elements react with non-metals?
Electrovalent compounds formed
122
Give examples of electrovalent compounds formed when group 1 elements react with non-metals.
Sodium chloride (NaCl), potassium bromide (KBr)
123
What is formed when group 1 elements react with hydrogen?
Ionic hydrides formed
124
Give examples of ionic hydrides formed when alkali metals react with hydrogen.
Lithium Hydride (LiH), Sodium hydride (NaH)
125
Name the elements in group 17 [VII A]
Fluorine, chlorine, bromine, iodine, astatine
126
What is the valency of elements in group 17 [VII A]?
Univalent [7 valence electrons]
127
Give 3 points about the nature of elements in group 17
1. Highly reactive
2. Highly electronegative
3. Non-metals
(non-metallic in nature)
128
Which elements in group 17 are gaseous?
Fluorine and Chlorine
129
Which elements in group 17 are liquid?
Bromine
130
Which elements in group 17 are solid at room temp?
Iodine
131
What kind of conductivity do elements of group 17 have?
Bad or non-conductors of heat and electricity
132
What is the nature of group 17 elements with reference to reduction/oxidization?
They are strong oxidizing agents.
133
Why are the elements of group 17 strong oxidizing agents?
This is because they are electron acceptors as they have 7 valence electrons and need only 1 to attain stable electronic configuration.
134
What is the nature of group 17 elements with reference to electronegativity?
They have high electronegativity. Electronegative character decreases from Fluorine to Iodine.
135
What is formed when group 17 elements react with non-metals?
Covalent compounds are formed.
136
Give examples of covalent compounds formed when group 17 elements react with non-metals.
Hydrochloric acid (HCl), Phosphorus trichloride (PCl₃), Disulfur dichloride (S₂Cl₂)
137
What is formed when group 17 elements react with hydrogen?
Covalent hydrides are formed when group 17 elements react with hydrogen.
138
Give examples of covalent hydrides formed when group 17 elements react with hydrogen.
Hydrogen Fluoride (HF), Hydrochloric acid (HCl), Hydrogen Bromide (HBr), Hydrogen Iodide (HI)
139
What does periodicity in properties of elements mean?
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
When do definite intervals occur in the MPT?
The definite intervals are after difference of either 2 or 8 or 18 or 32 in atomic numbers.
141
What are periodic properties?
The properties which appear at regular intervals in the periodic table are called periodic properties.
142
Name the periodic properties found in the MPT.
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
What is the reason for periodicity in properties in periods & groups?
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
What does atomic radius mean?
It is the distance between the centre of the nucleus and the outermost shell of the atom.
145
What does ionisation potential mean?
It is the amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom.
146
What does electron affinity mean?
It is the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.
147
What does electronegativity mean?
It is the tendency of an atom to attract electrons to itself when combined in a compound
148
What does metallic character mean?
In terms of electron loss or gain, an element is a metal if it gains one or more electrons
149
What does non-metallic character mean?
In terms of electron loss or gain, an element is a non-metal if it loses one or more electrons
150
What is the unit of atomic radii?
Angstron unit = Å
151
What is bond length?
Bond length or bond distance is defined as the average distance between nuclei of two bonded atoms in a molecule.
152
Which factors affect the atomic size?
1. Number of shells
| 2. Nuclear charge
153
What effect does no. of shells have on atomic size?
Atomic size increases if no. of shells increases
154
Why does atomic size increase on increase of no. of shells?
This is because when no. of shells increases, the distance of the outermost shell from the nucleus increases.
155
What effect does nuclear charge have on atomic size?
Atomic size decreases when nuclear charge increases.
156
Why does atomic size decrease with increase in nuclear charge?
When nuclear charge increases, the electrons in the outermost shell are attracted with increasing force.
157
What is nuclear charge?
It is the positive charge on the nucleus of an atom.
158
What is nuclear charge equivalent to?
It is equivalent to the atomic no. of an atom.
159
What is the trend in atomic size across a period?
Atomic size decreases across a period left to right.
160
What remains the same across a period?
The no. of shells remains the same, and therefore the atomic size remains unaffected
161
What increases across a period?
Nuclear charge increases across a period, and therefore, the atomic size decreases.
162
Which elements have the largest and smallest atomic radius in period 2?
Lithium has the largest atomic radius and Fluorine has the smallest atomic radius.
163
Why does Neon have a larger atomic radius despite being the last in the period?
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
How is a cation formed?
It is formed by loss of electron. Na - 1e- --> Na 1+
165
Why is a cation smaller than the parent neutral atom?
The remaining electrons in the cation are strongly attracted by the nucleus, thus decreasing the cation size.
166
What is the trend in atomic size down a group?
Atomic size (radii) increases down a group
167
What increases down a group?
No. of shells increases down a group, therefore atomic size increases.
168
Why does no. of shells increase down a group?
New shells are added with increasing atomic no.
169
Why does atomic size increase on increase in no. of shells down a group?
Nuclear distance from valence electrons increases on increase in no. of shells down a group
170
What happens to nuclear charge down a group?
Nuclear charge increases down a group
171
Why does nuclear charge increase down a group?
Nuclear charge is equivalent to the atomic no. of an element, and it increases as atomic no. increases down a group
172
What happens to at. size as nuclear charge increases down a group?
Atomic size should decrease as nuclear charge increases down a group.
173
Why does overall atomic radius increases down a subgroup?
Increase in no. of shells dominates over increase in nuclear charge, therefore overall atomic radius increases.
174
What is the first ionisation potential?
Energy required to remove 1st electron is called first I.P.
| M (atom) -> M+ (ion) + e- (electron)
175
What is the second ionisation potential?
Energy required to remove second electron is called second I.P.
M+ -> M++ + e-
176
Which I.P. is more?
The second I.P. is more.
177
What is the unit of I.P.?
Electron volt = e V
178
Name the factors which affect the I.P.
1. Atomic Size
| 2. Nuclear charge
179
How does atomic size affect ionisation potential?
When atomic size increases, I.P. decreases.
180
Why does I.P. decrease when atomic size increases?
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
How does nuclear charge affect I.P.?
When nuclear charge increases, I.P. increases.
182
Why does I.P. increase when nuclear charge increases?
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
What is the trend in I.P. across a period from left to right?
I.P. increases across a period from left to right.
184
How do the factors affecting I.P. vary across a period?
The atomic radii decreases, so I.P. increases.
| The nuclear charge increases, so I.P. increases.
185
Which element has the highest I.P.?
The element helium [He] has the highest I.P.
186
Which element has the lowest I.P.?
The element caesium [Cs] has the lowest I.P.
187
Why do metals have low I.P. as compared to non-metals?
Metals lose electrons and thus, have a lower I.P. than non-metals.
188
Why does francium not have the lowest I.P. despite being in the last period?
Francium [Fr] is radioactive.
189
What is the trend in I.P. down a group?
I.P. decreases down a group
190
Why does overall I.P. decrease down a group?
Increase in atomic radii dominates over increase in nuclear charge. Therefore, overall I.P. decreases down a group.
191
How do the factors affecting I.P. vary down a group?
Atomic radii increases down a group -> I.P. decreases
| Nuclear charge increases -> I.P. should increase
192
What is the 1st electron affinity?
The energy liberated when an atom in the gaseous state accepts an electron to form an anion.
193
What is the unit of electron affinity?
Electron volt = e V
194
With which sign is electron affinity represented?
Electron affinity is represented by a negative sign
195
Name the factors which affect electron affinity
1. Atomic size
| 2. Nuclear charge
196
How does atomic size affect electron affinity?
When atomic size increases, electron affinity decreases.
197
Why does electron affinity decrease on increase of atomic size?
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
How does nuclear charge affect electron affinity?
When nuclear charge increases, electron affinity increases.
199
Why does e.a. increase on increase of nuclear charge?
Increase in nuclear charge increases the tendency of an atom to accept electrons.
200
What is the trend in electron affinity across a period with relation to a.r and n.c?
Atomic radii decreases - Electron affinity increases
| Nuclear charge increases - Electron affinity increases
201
What is the e.a highest and lowest for?
Highest for halogens [group 17]
| Least for alkali metals [group 1]
202
Why does neon have e.a zero?
Inert gases with stable electronic configuration find it difficult to accept electrons.
203
Why do inert gases not form ions?
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
What is the trend in e.a across a period?
E.A. increases across a period from left to right
205
What is the trend in electron affinity down a group with relation to a.r and n.c?
Atomic radii increases - Electron affinity decreases
| Nuclear charge increases - Electron affinity should increase
206
Why does e.a. decrease down a group?
Increase in atomic radii dominates over increase in nuclear charge. Hence, overall e.a. decreases
207
What is the trend in electron affinity down a group?
E.A. decreases down a group
208
What is the effect of E.A. on the electronegative/oxidizing nature of an element?
Greater the value of E.A., more electronegative, or more oxidizing is the element
209
Which element has greater oxidization potential - electropositive or electronegative?
More electronegative element
210
Define electronegativity
The tendency of an atom to attract electrons to itself when combined in a compound
211
How does electronegativity affect the nature of bond?
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
Name the factors which affect E.N.?
1. Atomic size
| 2. Nuclear charge
213
How does atomic size affect E.N.?
When atomic size increases, E.N. decreases.
214
Why does E.N. decrease on increase of atomic size?
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
How does nuclear charge affect E.N.?
Nuclear charge increases - E.N. increases
216
Why does E.N. increase on increase of nuclear charge?
When nuclear charge increases, the tendency of an atom to accept electrons increases
217
Name the factors affecting E.N.
1. Atomic size
| 2. Nuclear charge
218
How do the factors affecting E.N. vary across a period?
Atomic radii decreases -> E.N. increases
| Nuclear charge increases -> E.N. increases
219
Which elements have high electronegativity?
Non-metallic elements
220
What is the trend in non-metallic character across a period?
It increases from left to right across a period
221
Which is the most E.N. element in the periodic table?
Fluorine
222
Which is the least E.N. element in the periodic table?
Caesium
223
Who do noble gases have zero E.N.?
They have a complete octet and hence do not attract electrons to themselves
224
What is the trend in overall E.N. across a period?
It increases across a period - left to right
225
Which elements is E.N. highest for?
Halogens
226
How do the factors affecting E.N. vary down a group?
Atomic radii increases -> E.N. decreases
| Nuclear charge increases -> E.N. should increase
227
Why does overall E.N. decrease down a group?
Increase in atomic radii dominates over increase in nuclear charge. Therefore, overall E.N. decreases
228
Define metallic character
An atom is said to be a metal if it loses one or more electrons when supplied with energy
229
Define non-metallic character
An atom is said to be a non-metal if it gains one or more electrons when supplied with energy
230
Name the factors influencing M.C/N.M.C
1. Atomic radii
| 2. Ionisation potential
231
How does atomic radii influence M.C/N.M.C?
A.R. increases - Metallic character increases
| A.R. increases - Non-metallic character decreases
232
How does I.P. influence M.C/N.M.C?
I.P. increases - Metallic character decreases
| I.P. increases - Non-metallic character increases
233
Give characteristics of metallic atoms
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
Give characteristics of non-metallic atoms
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
How does the reactivity of an element depend on its tendency to lose or gain electrons?
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
Are metals good reducing or oxidizing agents?
Metals are good reducing agents
237
Are non-metals good reducing or oxidizing agents?
Non-metals are good oxidizing agents
238
Give the trend in character across a period
Metallic character decreases across a period
| Non-metallic character increases across a period
239
How do the factors affecting character vary across a period?
A.R. decreases -> MC decreases, NMC increases
| I.P. increases -> MC decreases, NMC increases
240
How do the factors affecting character vary down a group?
A.R. increases -> MC increases, NMC decreases
| I.P. decreases -> MC increases, NMC decreases
241
Give the characteristics of the elements at the bottom of a group
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
Why do elements at the bottom of a group form ions from metals most readily/ are most reactive?
Thye have a large atomic size, the lowest I.P. and thus, electrons are loosely held
243
Give the trend in character down a group
Metallic character increases down a group
| Non-metallic character decreases down a group
244
Give the formula for a metal atom losing an electron
M -> M+ (ion) + e- (electron)
245
Give the formula for a non-metal atom gaining an electron
N + e- -> N- (ion)
246
What are light metals?
Elements arranged in the periodic table having a neutron/proton ratio around 1 are stable elements
247
Which light metals are stable elements?
light metals Na and K
248
What are heavy metals?
Elements with an n/p ratio above 1.5 are considered radioactive unstable elements
249
Give an example of a heavy metal that is a radioactive unstable element
Uranium
250
Give the steps to calculating if an element is stable or unstable
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
