Periodicity of Valence or Oxidation States of Elements

Periodicity of Valence or Oxidation States of Elements - Definition, Examples, Types, FAQs

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Valency Definition: The valencies of all 118 elements of an atom refers to the number of electrons it gains or losses to complete its outermost shell. By completing their outer shell, or octet, atoms become more stable (8 electrons in the outermost shell). As a result, periodicity of valence or oxidation states of elements interacts with other atoms or participate in chemical reactions in which they can give or lose electrons to other atoms or take or acquire electrons from other atoms to develop a stable state.

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Valency Concept:
What are valence electrons, and what do they do?
How do you determine an element's valency?
Oxidation Number vs Valency
Types of Valency
Electrovalence
Covalency
The Value of the First 30 Elements

As a result, they tend to interact with other atoms or participate in chemical reactions in which they give or lose electrons to other atoms or accept or gain electrons from other atoms to complete their octet and establish a stable state. So, the reason for bond formation or a chemical reaction between atoms is that atoms are always fighting to establish a stable state by clinging to each other. When an atom has one shell, for example, the stable state is obtained when it possesses two electrons. When an atom has two or more shells or orbits, the stable state is reached when the outermost shell possesses eight electrons. Hydrogen, for example, contains only one electron in the outermost or valence shell of its atom.

To complete electrons in its outermost orbit, it must add one more electron to its outermost shell or orbit. Similarly, the question arises what the valency of oxygen electronic arrangement of oxygen valence electrons is 2, 6. In its outermost shell or orbit, it has 6 electrons. As a result, it will require two additional electrons to complete the octet. Magnesium, with an atomic number of 12, has an electrical configuration of 2, 8, 2. To complete the octet, it will require 6 electrons. However, giving Mg 6 electrons is a difficult task for an atom. As a result, Magnesium's octet will be completed by donating or giving its two valence electrons to atoms of other elements.

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Valency Concept:

Noble gases are the least reactive because their periodicity of valence or oxidation states of elements shell is totally filled. The capacity of other elements to achieve the noble gas configuration determines their reactivity. For example: neon (Outermost shell is totally filled).

Electron_shell_010_Neon

What are valence electrons, and what do they do?

The valence electrons are the total number of electrons present in an atom's outermost orbit or energy shell. Magnesium's electrical configuration, for example, is 2, 8, 2. It means that the K orbit contains two electrons, the L orbit has eight, and the M orbit has two. In Magnesium's outermost orbit, there are two electrons. As a result, it contains two valence electrons.

How do you determine an element's valency?

The loss or gain of electrons relates to valency, which is distinct from the total amount of electrons in an atom. Since, Sodium valency has an atomic number of 11, it contains 11 electrons but a valency of 1. It is easier for it to lose 1 electron than to add 7 electrons to reach stability of periodicity of valence or oxidation states of elements of complete its octet since its electronic distribution is (2, 8, 1).

As a result, its valency is 1. Similarly, the valency of oxygen and calcium valency with atomic number 8 and electronic configuration (2, 6) is 2 since gaining two electrons to complete its octet is easy for it. On the other hand, atoms having the electrical configuration (2, 7) tend to gain one electron rather than lose all their seven electrons since gaining one electron is easier than losing seven. As a result, its valency is 1.

There are also elements in the periodicity of valence or oxidation states of elements with zero valency, such as noble gases like helium (He), neon (Ne), and argon (Ar), because their outermost shell has 8 electrons, and their octet is complete. As a result, they do not react with other atoms or elements, or they are believed to be the least reactive, and they are therefore referred to as inert gases.

Furthermore, valency affects the charge on an atom when it forms an ion after losing or acquiring electrons. The charge on a sodium ion, for example, is -1 because it loses one electron in a process to achieve a perfect electronic configuration (2, 8) or a filled outermost shell or orbit. Similarly, when Magnesium loses two electrons, the charge of the ion is +2 (Mg²⁺). Because of its electrical structure (2, 8, 2), it has a tendency to lose two electrons.

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Oxidation Number vs Valency

Both oxidation number and valency are expressions that refer to an atom's valence electrons. The main difference between the valency along with oxidation number is valence which in reference to the max-to-max number of electron which an atom can lose/gain/share fulfil the stability, meanwhile oxidation number is maximum number of electron which an atom can lose/gain for formation of the bond with another atom. The elements with valency 1 are diatomic molecules.

Types of Valency

Atoms unite to complete their octet, resulting in the production of compounds that can be either ionic or covalent. As a result, there are two forms of valency of all elements, as stated below.

Electrovalence

Electrovalence is a property of atoms that produce ionic or electrovalent compounds. The number of electrons lost or gained by an atom to establish a stable state or complete its octet is referred to as electrovalence. Positive ions (cations) are formed when an atom loses electrons, and their valency is called positive electrovalence. Negative ions (anions) are formed when atoms gain electrons, and their valency is called negative electrovalence.

So, electrovalence refers to the valency of ionic (electrovalent) compounds or electrovalent compounds that are generated when metals and non-metals come together. For example, sodium chloride (NaCl) is a non-metal with a metal component. The transfer of electrons from Sodium to Chlorine forms the chemical link between the atoms of sodium and chlorine in the Sodium Chloride molecule. Chlorine gains one electron whereas sodium loses one. As illustrated below, the electrovalence of sodium and chlorine is 1 because Na loses one electron and Cl gains one.

Na⁺+Cl^-→NaCl

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Covalency

Covalency Definition: The number of electrons shared by atoms or elements during the formation of a covalent molecule is referred to as the number of electrons shared by the atoms or elements. The chemical link between non-metals is established in covalent compounds for periodicity of valence or oxidation states of elements. Covalent compounds are formed when atoms of one non-metal unite with atoms of other non-metals to create molecules. The atoms do neither gain nor lose electrons in this situation; instead, they share electrons. In Methane (CH₄), for example, one carbon atom, which requires four electrons to complete its octet, unites with four hydrogen atoms, each of which requires one electron to complete the two electrons necessary in its initial orbit.

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The Value of the First 30 Elements

Consider the valency chart or valency of the first 30 elements in the periodic table with valency.

VALENCY TABLE
ELEMENT NAME	VALENCY OF ELEMENTS	ELEMENT’S ATOMIC NUMBER
VALENCY OF HYDROGEN	1	1
VALENCY OF HELIUM	0	2
VALENCY OF LITHIUM	1	3
VALENCY OF BERRYLLIUM	2	4
VALENCY OF BORON	3	5
VALENCY OF CARBON	4	6
VALENCY OF NITROGEN	3	7
VALENCY OF OXYGEN	2	8
VALENCY OF FLORINE	1	9
VALENCY OF NEON	0	10
VALENCY OF SODIUM	1	11
VALENCY OF MAGNESIUM	2	12
VALENCY OF ALUMINIUM	3	13
VALENCY OF SILICON	4	14
VALENCY OF PHOSPHOROUS	3	15
VALENCY OF SULPHUR	2	16
VALENCY OF CHLORINE	1	17
VALENCY OF ARGON	0	18
VALENCY OF POTASSIUM	1	19
VALENCY OF CALCIUM	2	20
VALENCY OF SCANDIUM	3	21
VALENCY OF TITANIUM	4	22
VALENCY OF VANADIUM	5,4	23
VALENCY OF CHROMIUM	2	24
VALENCY OF MANGANESE	7,4,2	25
VALENCY OF IRON	2,3	26
VALENCY OF COBALT	3,2	27
VALENCY OF NICKEL	2	28
VALENCY OF COOPER	2,1	29
VALENCY OF ZINC	2	30

The valency of first 20 elements is given in above table.

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Frequently Asked Questions (FAQs)

1. What is variable valency, and how does it work?

Variable valency is a term used to describe items that display many types of valency. In different compounds, these elements have varying valences. Transition elements such as iron valency, mercury, and copper, for example, have different valences. In FeSO₄, iron valency of 2, whereas in FeCl₃, it has a valency of 3. Copper is divided into two valences: 1 and 2. Mercury has two different valences (1 and 2) as well.

2. Is it possible for Valency to be negative?

Atom charge is the loss or gain of an electron. Positive charge is obtained by donating an electron, while negative charge is obtained by receiving an electron. Charge has both positive and negative signals, whereas valency has no sign.

3. What is the valency of copper?

The valency of copper (Cu) is 2,1.

4. What is the valency of Manganese?

The valency of Manganese is 7,4,2.

5. Write the valency of mercury, nickel, and cobalt.

Valency of mercury = 1.

Valency of nickel = 3,2.

Valency of cobalt = 2.

6. Why do some elements, like carbon, have a wide range of oxidation states?

Elements like carbon can form multiple covalent bonds and participate in various types of bonding arrangements. This flexibility allows them to adopt a wide range of oxidation states, from -4 to +4 in carbon's case.

7. Why do some elements, like phosphorus, have oxidation states that don't follow the octet rule?

Elements like phosphorus can form expanded octets, allowing them to have more than eight electrons in their valence shell. This ability enables phosphorus to exhibit oxidation states higher than +3, such as +5, which doesn't follow the simple octet rule.

8. How do the oxidation states of p-block elements compare to those of s-block elements?

P-block elements generally exhibit a wider range of oxidation states compared to s-block elements. While s-block elements typically have a single common oxidation state (+1 or +2), p-block elements can have multiple oxidation states due to their ability to form covalent bonds and share electrons.

9. How do the oxidation states of transition metals change across a period?

Across a period in the transition metals, the maximum oxidation state generally increases from left to right. This trend is due to the increasing nuclear charge and the availability of more electrons for bonding or removal.

10. What role do lone pairs play in determining the oxidation states of elements?

Lone pairs of electrons can influence an element's oxidation states by affecting its ability to form bonds or participate in reactions. Elements with lone pairs may have a wider range of oxidation states as these electrons can be involved in bonding or remain non-bonding.

11. How does the concept of electronegativity difference help predict the oxidation states in ionic compounds?

The electronegativity difference between elements in a compound can indicate the likely direction of electron transfer. The more electronegative element typically gains electrons and has a negative oxidation state, while the less electronegative element loses electrons and has a positive oxidation state.

12. How does electronegativity affect the oxidation states of elements?

Electronegativity influences oxidation states by determining how atoms share or transfer electrons in chemical bonds. More electronegative elements tend to have negative oxidation states, while less electronegative elements often have positive oxidation states.

13. How does the lanthanide contraction affect the oxidation states of f-block elements?

The lanthanide contraction leads to a decrease in atomic and ionic radii across the lanthanide series. This can influence the stability of certain oxidation states, often favoring the +3 state for many lanthanides and actinides.

14. How does the concept of effective nuclear charge relate to the periodicity of oxidation states?

Effective nuclear charge increases across a period, making it harder for elements on the right side of the periodic table to lose electrons. This trend contributes to the periodicity of oxidation states, with elements on the right typically having more negative oxidation states.

15. How does the concept of isoelectronic species relate to oxidation states?

Isoelectronic species have the same number of electrons but different nuclear charges. This concept can help explain why some ions with the same oxidation state have similar properties, even if they come from different elements.

16. How does the octet rule influence the valence states of elements?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. This principle often determines the most common valence states for many elements, especially those in the main groups.

17. What is the relationship between an element's common oxidation states and its electron configuration?

An element's common oxidation states are often related to its electron configuration, particularly the number of valence electrons. Elements tend to lose, gain, or share electrons to achieve a stable electron configuration, which determines their most common oxidation states.

18. Why do noble gases typically have a valence state of zero?

Noble gases have a complete outer shell of electrons, making them extremely stable. As a result, they rarely form chemical bonds or participate in reactions, leading to a valence state of zero in most cases.

19. How do the oxidation states of elements in the same group vary in their oxides?

Elements in the same group often form oxides with similar formulas, but the stability of these oxides can vary. Generally, the highest oxidation state becomes more stable for heavier elements in a group, reflecting the increased ability to accommodate multiple bonds.

20. Why do some elements, like nitrogen, have oxidation states that seem to "skip" certain values?

Elements like nitrogen may have oxidation states that appear to skip certain values due to the stability of particular electronic configurations or molecular structures. For nitrogen, the common oxidation states are -3, +3, and +5, with +1, +2, and +4 being less common or unstable in many compounds.

21. What is the difference between valence state and oxidation state?

Valence state typically refers to the number of electrons an atom can lose, gain, or share in chemical bonding. Oxidation state, on the other hand, represents the degree of oxidation of an atom in a chemical compound, which may or may not be equal to its valence state.

22. How do the oxidation states of main group elements typically change across a period?

Across a period, the maximum oxidation state of main group elements generally increases from left to right. This trend is due to the increasing number of valence electrons and the element's ability to lose or share more electrons.

23. How do the oxidation states of halogens vary across the group?

Halogens can exhibit multiple oxidation states, ranging from -1 to +7. The -1 state is most common, but higher oxidation states become more stable as you move down the group due to the increasing stability of d-orbitals.

24. Why do some elements, like nitrogen, have both positive and negative oxidation states?

Elements like nitrogen can both gain and lose electrons in chemical reactions. Nitrogen's ability to form multiple covalent bonds and its electronegativity allow it to exhibit oxidation states ranging from -3 to +5, depending on the compound.

25. What is disproportionation, and how does it relate to oxidation states?

Disproportionation is a reaction where a single element in one oxidation state is simultaneously oxidized and reduced to form two different oxidation states. This process demonstrates how some elements can exist in multiple oxidation states.

26. What is the periodicity of valence states?

The periodicity of valence states refers to the recurring pattern of oxidation states that elements exhibit across periods and groups in the periodic table. This pattern is closely related to the electronic configuration of elements and their position in the periodic table.

27. How does the valence state of an element relate to its position in the periodic table?

An element's valence state is directly related to its position in the periodic table. Generally, elements in the same group tend to have similar valence states, while the maximum valence state often increases across a period from left to right.

28. What role do unpaired electrons play in determining the oxidation states of transition metals?

Unpaired electrons in the d-orbitals of transition metals can participate in bonding, allowing these elements to form compounds with various oxidation states. The number of unpaired electrons often correlates with the range of possible oxidation states.

29. Why do transition elements often have multiple oxidation states?

Transition elements often have multiple oxidation states because they have partially filled d-orbitals. These d-electrons can be easily removed or shared, allowing the element to form compounds with different oxidation states.

30. What is the relationship between group number and maximum oxidation state for main group elements?

For main group elements, the maximum oxidation state often corresponds to the group number. For example, elements in group 1 typically have a maximum oxidation state of +1, while elements in group 7 can have a maximum oxidation state of +7.

31. How does the inert pair effect influence the oxidation states of heavier p-block elements?

The inert pair effect causes the two s electrons in the outermost shell of heavier p-block elements to become less reactive. This can lead to a preference for oxidation states that are two less than the group number, especially in groups 13-15.

32. Why do some elements, like copper and gold, have unusual oxidation states compared to others in their group?

Elements like copper and gold can have unusual oxidation states due to factors such as the stability of certain electron configurations (e.g., d10) and relativistic effects. These factors can make some oxidation states more favorable than what would be expected based on group trends.

33. Why do some elements, like sulfur, have a wider range of oxidation states than others in the same group?

Elements like sulfur can form expanded octets, allowing them to accommodate more than eight electrons in their valence shell. This ability, combined with the availability of d-orbitals, enables sulfur to exhibit a wider range of oxidation states compared to oxygen, for example.

34. How does the stability of half-filled and fully-filled d-orbitals affect the oxidation states of transition metals?

Half-filled (d5) and fully-filled (d10) d-orbital configurations are often more stable. This can lead to a preference for oxidation states that result in these electronic configurations, influencing the chemistry of transition metals.

35. How do the oxidation states of lanthanides differ from those of transition metals?

Lanthanides typically show less variation in oxidation states compared to transition metals. The +3 oxidation state is most common for lanthanides, while transition metals often exhibit a wider range of oxidation states due to their partially filled d-orbitals.

36. What is the relationship between an element's oxidation states and its ability to catalyze chemical reactions?

Elements with multiple accessible oxidation states, particularly transition metals, often make good catalysts. The ability to easily switch between oxidation states allows these elements to facilitate electron transfer in reactions, making them effective in catalytic processes.

37. Why do some compounds, like metal carbonyls, have oxidation states that seem counterintuitive?

In metal carbonyls, the metal often has a low or even zero oxidation state despite being bonded to several oxygen atoms. This is due to the strong π-backbonding between the metal and CO ligands, which effectively neutralizes the charge on the metal.

38. How does the stability of different oxidation states vary for elements across a group?

As you move down a group, higher oxidation states often become more stable due to the increasing availability of d-orbitals and the decreasing energy gap between s and p orbitals. This trend is particularly noticeable in groups 13-17.

39. What is the relationship between an element's oxidation states and its ability to form coordination compounds?

Elements that can exist in multiple oxidation states, particularly transition metals, are often good at forming coordination compounds. The availability of different oxidation states allows these elements to act as versatile central atoms in complex ions.

40. Why do some transition metals prefer certain oxidation states over others?

The preference for certain oxidation states in transition metals is often related to the stability of specific electron configurations. States that result in half-filled or fully-filled d-orbitals (d5 or d10) are often more stable due to their symmetric electron distribution.

41. What is the significance of the +2 oxidation state in transition metals?

The +2 oxidation state is common in many transition metals because it often results from the loss of the outermost s electrons, leaving the d-orbitals intact. This state can be particularly stable for elements with nearly half-filled or fully-filled d-orbitals.

42. How do the oxidation states of actinides compare to those of lanthanides?

Actinides generally show a wider range of oxidation states compared to lanthanides. While lanthanides predominantly exist in the +3 state, actinides can exhibit multiple oxidation states due to the involvement of 5f, 6d, and 7s orbitals in bonding.

43. Why do some elements, like iron, have certain oxidation states that are more common in biological systems?

In biological systems, the availability of certain oxidation states is often related to their redox potential and stability under physiological conditions. For iron, the +2 and +3 states are common because they are easily interconvertible and suitable for electron transfer processes in proteins.

44. How does the concept of electrochemical series relate to the stability of different oxidation states?

The electrochemical series ranks elements based on their standard reduction potentials. This ranking can help predict the stability of different oxidation states, with more positive reduction potentials generally indicating more stable reduced forms (lower oxidation states).

45. What is the relationship between an element's oxidation states and its position in the reactivity series?

Elements higher in the reactivity series tend to form more stable compounds in their higher oxidation states, while elements lower in the series may prefer lower oxidation states. This trend is related to the element's ability to lose electrons and form stable ionic compounds.

46. How do the oxidation states of metalloids differ from those of metals and non-metals?

Metalloids often exhibit a range of oxidation states that reflect their intermediate character between metals and non-metals. They can form both positive and negative oxidation states, depending on the electronegativity of the elements they're bonded to.

47. Why do some elements, like chlorine, have fractional oxidation states in certain compounds?

Fractional oxidation states occur in compounds where the oxidation state is an average over equivalent atoms. For example, in the peroxide ion (O2^2-), each oxygen atom has an oxidation state of -1, which is not a whole number but represents the average charge per atom.

48. How does the concept of formal charge relate to oxidation states?

Formal charge and oxidation state are both methods of assigning electron ownership in compounds, but they use different rules. While formal charge considers the number of valence electrons an atom "owns" in a molecule, oxidation state is based on the hypothetical complete transfer of electrons to the more electronegative atom.

49. What role do oxidation states play in understanding redox reactions?

Oxidation states are crucial in understanding redox reactions as they help identify which species are being oxidized (increase in oxidation state) and which are being reduced (decrease in oxidation state). This concept is fundamental to balancing redox equations and predicting reaction outcomes.

50. How does the concept of electron affinity relate to the tendency of elements to form negative oxidation states?

Electron affinity, the energy change when an atom gains an electron, is closely related to an element's tendency to form negative oxidation states. Elements with high electron affinities, like halogens, are more likely to form stable negative ions and thus have negative oxidation states in many compounds.

51. How do the oxidation states of elements change in different types of chemical bonds (ionic, covalent, metallic)?

In ionic bonds, elements often adopt their highest (for metals) or lowest (for non-metals) oxidation states. In covalent bonds, oxidation states can be more variable and are determined by electron-sharing. In metallic bonds, atoms typically have low or zero oxidation states due to delocalized electrons.

52. How does the concept of electronegativity equalization relate to the oxidation states in a molecule?

Electronegativity equalization suggests that in a molecule, electrons are distributed to minimize electronegativity differences. This concept can help explain why some atoms in a molecule may have oxidation states that differ from their typical values in simpler compounds.

53. What is the significance of the +3 oxidation state in group 13 elements?

The +3 oxidation state is significant for group 13 elements as it represents the loss of all three valence electrons. However, due to the inert pair effect, heavier elements in this group (like thallium) may prefer the +1 state, showing the importance of considering multiple factors when predicting oxidation states.

54. How do the oxidation states of elements relate to their ability to form complex ions?

Elements that can exist in multiple oxidation states, especially transition metals, are often good at forming complex ions. The availability of different oxidation states allows these elements to act as versatile central atoms, accepting various numbers and types of ligands to form stable complexes.

55. Why is it important to consider both the highest and lowest common oxidation states of an element?

Considering both the highest and lowest common oxidation states provides a complete picture of an element's chemical behavior. The highest oxidation state often reflects the group number and shows the maximum number of electrons an atom can lose or share. The lowest state indicates the element's reducing power and its ability to gain electrons. Together, these extremes help predict the element's reactivity and the types of compounds it can form.