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Inorganic Chemistry > Classification Of Elements And Periodicity In Elements > Flashcards

Classification Of Elements And Periodicity In Elements Flashcards

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1
Q

State Mendeleev ‘s Periodic law.

A

The properties of the elements are periodic function of their atomic weight.

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2
Q

Moseley plotted graph between which two quantities, and stated that atomic no. is more fundamental property of an element.

A

Sq. root of Frequency and atomic no. which have a straight line.

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3
Q

State modern periodic law?

A

The physical and chemical properties of the elements are periodic functions of their atomic numbers.

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4
Q

Tell word root (0-9) for IUPAC notation of elements having Z>100?

A

0 - nil
1 - un
2 - bi
3 - tri
4 - quad
5 - pent
6 - hex
7 - sept
8 - oct
9 - enn

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5
Q

Tell IUPAC name of Z= 196 and symbol also.

A

Name:- Unennhexium
Symbol:- Ueh

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6
Q

Which quantum no defines the main energy level, i.e shell.

A

Principal Quantum no.(n)

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7
Q

Why is He placed in the p-block with noble gases?

A

Because it has a completely filled valence shell (1s²) and as a result exhibits properties characteristics of other noble gases.

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8
Q

Why is hydrogen behave like Group 1 nd Group 17 elements?

A

Group 1:- Has only one s-electron.
Group 2:- Can gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 ( halogen family) elements.

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9
Q

Do S-block elements have higher ionisation enthalpy?

A

No, as they loose the outermost electron(s) readily to form 1+ ion( in the case of alkaline metals) or 2+ ion (in the case of alkaline earth metals.

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10
Q

Trend of metalic character and reactivity in S- block?

A

Both increase as we go down the group.

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11
Q

Which two elements in S-block elements don’t form predominantly ionic compounds.

A

Lithium and Beryllium

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12
Q

General electronic configuration of S-block and P- block?

A

S block:- ns¹ and ns²
P-block:- ns²np¹ to ns²np⁶

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13
Q

Why does noble gas exhibits very low chemical reactivity?

A

All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons.

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14
Q

Trend of metallic character in a group and that of non- mettal in a period?

A

Metalic character increases as we go up to down in a group, whereas non-metallic character increase from left to right in a period.

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15
Q

Why do Zn, Cd, Hg do not show most of the prop. of transition elements?

A

Because they have completely filled d-orbital, their electronic configuration is :- (n-1) d¹⁰ ns².

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16
Q

State General electronic configuration of d-block elements?

A

Electronic configuration:- (n-1)d¹–¹⁰ ns⁰–²

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17
Q

Starting and ending elements of Lanthanoids and Actinoids?

A

Lanthanoids:- Ce (Z=58) - Lu (Z=71)
Actinoids:- Th ( Z =90) - Lr (Z=103)

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18
Q

What are Transuranium elements?

A

The elements after uranium are Transuranium metals.

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19
Q

Name three metals which have low melting points?

A

Mercury as an exception is liq. at room temp.
Gallium (303K) and Caesium(303K) also have very low melting points.

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20
Q

Name two metals which have high melting and boiling points?

A

Boron and Carbon are exceptions.

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21
Q

What are metalloids give example?

A

Elements which show properties that are characteristic of both metals and non-metals, these are called Semi-metals or Metalloids,
Ex:- silicon, germanium, arsenic, antimony and tellurium.

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22
Q

What is the concept behind calculating the covalent radius?

A

Refer to official inorganic notes and write the explanation here.

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23
Q

Why does atomic size generally decrease within a period?

A

Because within the period the outer electrons are in the same valence shell and the effective nuclear charge increase as the atomic number increases resulting in the increased attraction of electrons to the nucleus.

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24
Q

Why does atomic size generally increase as be designed a group, state two reasons?

A

Reason 1:- As we descend the groups, the principle quantum number (n) increases and the valency electrons are father from the nucleus.

Reason 2:- This happens because the inner energy levels are filled with electrons which serve to shield the outer electrons from the pull of the nucleus.

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25
Q
  1. Cations are formed by?
  2. Anions are formed by?
A
  1. Removal of electrons from an atom.
  2. Gain of an electron by an atom.
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26
Q

Why is cations smaller than its parent atom, and anion larger in size than its parent atom.

A

A cation is smaller than its parent atom because it has fever electrons while its nuclear charge remains the same.

The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge.

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27
Q

What are isoelectronic species? Give examples.

A

Atoms and ions which contain the same number of electrons are called isoelectronic species.
Ex:- O²–, F–, Na+ and Mg²+

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28
Q

Define ionization enthalpy?

A

It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state.

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29
Q

Define second ionization enthalpy?

A

The energy required to remove the second most loosely bound electron.

30
Q

Ionization enthalpy can be positive and negative also, (True or False)?

A

False, energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive.

31
Q

Give reason why is the order of ionization enthalpy is; 3>2>1, also which IE wil you consider if only IE is written in question?

A

The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom. In the same way the third ionization enthalpy will be higher than the second and so on.
The term “ionization enthalpy” if not qualified is taken is the first ionization enthalpy.

32
Q

State the trend of ionization enthalpy across the period and group?

A

Across a period ionization enthalpy increases from left to right.
Across a group ionization enthalpy decreases as we descend in a group.

33
Q

What do you mean by shielding or screening effect?

A

The repulsion provided by the inner core electrons to the valency electrons is called shielding or screening effect.

34
Q

When is shielding highly effective?

A

In general shielding is effective in the orbitals in the inner shells are completely filled.

35
Q

Why does nuclear charge outweighs the shielding across a period ?

A

Across a period successive electrons are added to orbitals in the same principle quantum number and the shielding of the nuclear charge by the inner core of electrons does not increase very much to compensate for the increased attraction of the electrons to the nucleus thus across a period increasing nuclear charge outweighs the shielding.

36
Q

Why does shielding outweighs the nuclear charge across a group?

A

Down a group the outermost electron being increasingly farther from the nucleus, there is an increased shielding of the nuclear charge by the electrons in the inner energy levels.

37
Q

Why is first ionization enthalpy of boron is less than first ionization enthalpy of beryllium even when boron has higher nuclear charge?

A

Because an s-electron is attracted to the nucleus more than a p-electron. In beryllium the electron removed the during the ionization is an s-electron whereas the electron removed during ionization of boron is a p-electron. The penetration of a 2S-electron to the nucleus is more than that of a 2P-electron: hence the 2p-electron of boron is more shielded from the nucleus by the inner core of electrons then the 2s-electrons of beryllium.

38
Q

Why ionization enthalpy of oxygen is less than that of ionization enthalpy of nitrogen?

A

In nitrogen atom three 2P-electrons reside in different atomic orbitals(Hunds rule) whereas in the oxygen atom two of the four 2p-electrons must occupy the same 2p-orbital resulting in an increased electron-electron repulsion. consequently it is easier to remove the fourth 2P-electron from oxygen then it is to remove one of the three 2p- electrons from nitrogen.

39
Q

Define electron gain enthalpy?

A

When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the electron gain enthalpy (∆ēg H).

40
Q

Electron gain enthalpy is always -ve ? (True or False)

A

False, for many elements energy is released when an electron is added to the atom and the electron gain enthalpy is negative, however in the case of noble gases large positive electron gain enthalpy are observed because the electron has to enter the next higher principle quantum number leading to a very unstable electronic configuration.

41
Q

As atomic number increases electron gain enthalpy becomes more___________?

A

Negative.

42
Q

Trend of electron gain enthalpy across the period and a group?

A

Across a period electron gain enthalpy becomes more negative as we go from left to right.

Across a group electron gain enthalpy becomes less negative because the size of the atom increase in the at the electron would be further from the nucleus.

43
Q

Why is electron gain enthalpy of Oxygen and fluorine is less negative than that of the succeeding element.

A

This is because when an electron is added to oxygen or fluorine the added electron goes to the smaller n=2 Quantum level and suffer significant repulsion from the other electrons present in this level, for the n=3 Quantum level (S or Cl), the added electron occupies the larger region of space and the electron- electron repulsion is much less.

44
Q

Define electronegativity? Is it a measurable quantity? State the numerical scales of electronegativity?

A

A qualitative major of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity, it is not a measurable quantity.

Numerical scales of electronegativity of elements are Pauling scale mulliken Jeff scale, Allred-Rochow scale.

45
Q

Define electron affinity and its sign convention and its relation with electron gain enthalpy.

A

The negative of the enthalpy change for the process depicted:-
X(g) + e- ——> X– (g)

Is defined as electron affinity (Aè) it’s convention is quadratory to thermodynamic convention of chemistry.

∆eg H = -Aè - 5/2 RT

46
Q

Most electronegative element in periodic table is?

A

Fluorine

47
Q

The electronegativity of all elements is constant? (True or False)

A

False, electronegativity of any given element is not constant it varies depending on the element to which it is bound.

48
Q

State the trend of electronegativity across a period and a group?

A

Electronegativity generally increases across a period from left to right and decreases down a group.

49
Q

Relation between electron gain enthalpy ionization enthalpy electronegativity and atomic radii with each other?

A
  1. ∆eg H directly proportional to ionization enthalpy.
  2. ∆eg H directly proportional to electronegativity.
  3. ∆eg H inversely proportional to atomic radii.
50
Q

Non metallic elements have_______ tendency to _________electrons? ( Strong, beak, gain, loose)

A

Non metallic elements have strong tendency to gain electrons.

51
Q

Proportionality between electronegativity and metallic character and non metallic character also?

A

Electronegativity is directly proportional to non metallic characters and inversely proportional to metallic characters.

52
Q

The periodic trends of elements in the periodic table?

A

NCERT chemistry—>Part 1 —> class 11—> pg, 91.

53
Q

Valence of representative elements is usually equal to?

A

–> The no. Of electrons in their outermost orbitals and/or equal to eight minus the nom of outermost electrons as shown.

54
Q

Tell oxidation state of F, O and Na in OF2 and Na2O, also tell the analogy of +ve and -ve sign before no. in them.

A

NCERT class 11, part 1, Pg:- 92
Upper right sided para.

55
Q

Define oxidation state of an element in a particular compund?

A

Oxidation state of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.

56
Q

Predict formulas of compund which might be formed by pairs of elements;
a) silicon and bromine
b) aluminium and sulphur

A

SiBr4
Al2S3

57
Q

How can be use electronegativity to decide that who will take electron and who will donate.

A

The who is more EN- will take electron and the one who is less EN- will donate.

Ex:- In OF2, F takes e- and O donates e-, however in Na2O O takes e- and Na gives.

58
Q

How is Li and Be are unlike from their groups elements.

A

Lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals form compounds with pronounced covalent character, the other members of these groups predominantly form ionic compounds.

59
Q

Define Diagonal realationship and it’s cause, also behaviour of Li is similar to which element.

A

Similarities in properties of first group elements with the second elements of following group is called Diagonal relationship, it is caused due to same polarising power.

Polarising power= (charge)² /size

Li is similar in properties to Mg.

60
Q

State 6 reasons for anomalous properties of second period elements as compared to their other group members.

A
  1. Small size
  2. High electronegativity
  3. High charge/ radius ratio.
  4. Only 4 orbitals available for bonding i.e ( 2s and 2p), i.e maximum covalence of the first member of each group is 4.
  5. Cannot expand their valencies like second members of their groups.
  6. The first member of p-block elements displays greater ability to form pπ-pπ multiple bonds to itself and to other second period elements.
61
Q

State maximum covalency of first member of 2nd group.

A

4

62
Q

Which members of p block displays greater ability to form pπ-pπ bonds to itself and to other second period elements.

A

First member of p block.

63
Q

Are the oxidation state and covalency of Al in [AlCl(H2O)5]²+ same?

A

No, the oxidation state of Al is +3 and the covalency is 6.

64
Q

All chemical and physical properties are a manifestation of their electronic configuration of elements (T/F)?

A

True.

65
Q

Where is ionization energy is least and highest in periodic table.

A

Least on extreme left
Highest on extreme right

66
Q

Where is electronic gain enthalpy highest -ve and least -ve in periodic table.

A

Highest -ve in extreme right.
Least -ve in extreme left

67
Q

Why is Chemical reactivity high at extreme ends?

A

Due to least I.E on extreme left and highest Electron gain enthalpy on extreme right.

68
Q

Why does extreme left elements form most basic oxides and also why it’s vice versa is also true.

A

As extreme left elements can easily lose electron they form basic oxides whereas the extreme right elements have high tendency to gain electrons that why they form most acidic oxides.

69
Q

Give examples of amphoteric and neutral oxides?

A

Amphoteric Oxides:- Al2O3 , As2O3.
Neutral Oxides:- Co, NO, N2O.

70
Q

3d elements are more electro positive then group one and two elements (true or false)?

A

False.

71
Q

Why does the reverse trend of ionization energy, atomic size, electronegativity, electron affinity is generally observed in d block.

A

Following are the reasons:-
Lanthanoid contraction and poor shielding effect of inner shell electrons due to which instead of increasing of atomic size, atomic size decreases.

72
Q

Oxides of elements on right, left and centre are ____,_____,____respectively in nature.

A

Acidic, Basic, Amphoteric or Neutral.

Inorganic Chemistry (1 decks)

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