Periodic Classification of Elements

The periodic table is divided into rows and columns, where the rows are called periods, and the columns are called groups or families. Each element in the table is represented by its chemical symbol, which is a one- or two-letter abbreviation of its name. The elements in each row have the same number of electron shells, while the elements in each column have similar chemical properties.

The periodic table has helped scientists predict the properties of new elements and understand the behavior of known elements. The periodic table is also useful in predicting chemical reactions and the properties of compounds formed from different elements.

The periodic table is a fundamental tool in chemistry, and its discovery is considered one of the most significant achievements in science. The first version of the periodic table was developed independently by Dmitri Mendeleev and Julius Lothar Meyer in the 1860s. Since then, the periodic table has been revised and expanded as new elements have been discovered. Today, there are 118 elements in the periodic table, with the most recent additions being nihonium, moscovium, tennessine, and oganesson.

Döbereiner’s Triads were a set of three elements that were grouped together by the German chemist Johann Wolfgang Döbereiner in the early 1800s. He noticed that some groups of three elements had similar chemical properties, and that the atomic weight of the middle element was approximately equal to the average of the other two.

Döbereiner identified several such triads, including lithium, sodium, and potassium; calcium, strontium, and barium; and chlorine, bromine, and iodine. He called these groups “triads” and suggested that they represented a fundamental pattern in the properties of the elements.

The discovery of Döbereiner’s Triads was an important step in the development of the periodic table, as it suggested that there might be a regular pattern in the properties of the elements. However, Döbereiner’s Triads were ultimately found to be limited in their usefulness, as they only applied to a small number of elements and did not account for the full range of chemical properties exhibited by different elements.

Nonetheless, Döbereiner’s Triads helped lay the groundwork for later attempts to classify and organize the elements, and they remain an important part of the history of chemistry.

Here are some examples of Döbereiner’s Triads:

1. Calcium, Strontium, Barium

• Calcium (Ca) has an atomic weight of 40.08
• Strontium (Sr) has an atomic weight of 87.62
• Barium (Ba) has an atomic weight of 137.33
• The atomic weight of strontium is approximately the average of calcium and barium (i.e., (40.08 + 137.33)/2 = 88.70)

2. Chlorine, Bromine, Iodine

• Chlorine (Cl) has an atomic weight of 35.5
• Bromine (Br) has an atomic weight of 79.9
• Iodine (I) has an atomic weight of 126.9
• The atomic weight of bromine is approximately the average of chlorine and iodine (i.e., (35.5 + 126.9)/2 = 81.2)

3. Lithium, Sodium, Potassium

• Lithium (Li) has an atomic weight of 6.94
• Sodium (Na) has an atomic weight of 22.99
• Potassium (K) has an atomic weight of 39.10
• The atomic weight of sodium is approximately the average of lithium and potassium (i.e., (6.94 + 39.10)/2 = 23.02)

The Law of Octaves was proposed by the English chemist John Newlands in 1864. According to this law, the properties of the elements repeated every eighth element when they were arranged in order of increasing atomic weight.

Newlands observed that every eighth element had properties that were similar to the first element, much like the notes in an octave of music. He arranged the elements in rows of seven, and noted that the eighth element in each row had similar properties to the first element in the next row.

Newlands’ Law of Octaves was an early attempt to classify the elements, and it represented an important step towards the development of the periodic table. However, it was met with skepticism by many scientists of the time, who found the patterns to be too inconsistent to be useful. Additionally, Newlands’ approach did not account for the existence of elements that had not yet been discovered.

Despite these limitations, Newlands’ Law of Octaves played a role in shaping the later development of the periodic table, and it remains an important part of the history of chemistry.

Here are some examples of Newlands’ Law of Octaves:

1. Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, Fluorine

• Lithium (Li) has an atomic weight of 6.94
• Beryllium (Be) has an atomic weight of 9.01
• Boron (B) has an atomic weight of 10.81
• Carbon (C) has an atomic weight of 12.01
• Nitrogen (N) has an atomic weight of 14.01
• Oxygen (O) has an atomic weight of 16.00
• Fluorine (F) has an atomic weight of 19.00
• The eighth element in this series would be sodium (Na), which has an atomic weight of 22.99. Sodium has similar properties to lithium, the first element in the series.

2. Calcium, Strontium, Barium, Lanthanum, Cerium, Praseodymium, Neodymium

• Calcium (Ca) has an atomic weight of 40.08
• Strontium (Sr) has an atomic weight of 87.62
• Barium (Ba) has an atomic weight of 137.33
• Lanthanum (La) has an atomic weight of 138.91
• Cerium (Ce) has an atomic weight of 140.12
• Praseodymium (Pr) has an atomic weight of 140.91
• Neodymium (Nd) has an atomic weight of 144.24
• The eighth element in this series would be potassium (K), which has an atomic weight of 39.10. Potassium has similar properties to calcium, the first element in the series.

3. Iron, Cobalt, Nickel, Copper, Zinc, Gallium, Germanium

• Iron (Fe) has an atomic weight of 55.85
• Cobalt (Co) has an atomic weight of 58.93
• Nickel (Ni) has an atomic weight of 58.69
• Copper (Cu) has an atomic weight of 63.55
• Zinc (Zn) has an atomic weight of 65.38
• Gallium (Ga) has an atomic weight of 69.72
• Germanium (Ge) has an atomic weight of 72.63
• The eighth element in this series would be arsenic (As), which has an atomic weight of 74.92. Arsenic has similar properties to iron, the first element in the series.

Mendeleev’s Periodic Table is the first widely recognized and successful attempt to classify the elements. The Russian chemist Dmitri Mendeleev published his periodic table in 1869, and it quickly gained acceptance in the scientific community. The table was based on the observation that when the elements were arranged in order of increasing atomic weight, certain properties of the elements recurred at regular intervals.

Mendeleev’s table consisted of rows, or periods, and columns, or groups, of elements. Each element was assigned a unique atomic number, and elements with similar properties were grouped together in the same column. The elements were arranged in such a way that elements with similar properties fell into the same vertical column, and elements with similar electronic configurations fell into the same horizontal row.

Mendeleev’s table had several significant features, including:

1. Periodicity: Mendeleev arranged the elements in order of increasing atomic weight, and noted that the properties of the elements repeated at regular intervals. This periodicity allowed him to predict the properties of elements that had not yet been discovered.
2. Grouping by Valence: Mendeleev grouped elements by their valence, or the number of electrons in their outermost shell. This allowed him to predict the chemical properties of elements based on their position in the table.
3. Left Gaps for Unknown Elements: Mendeleev left gaps in his periodic table for elements that had not yet been discovered. By examining the properties of the elements in the same row as the gap, he was able to make predictions about the properties of the missing element.
4. Use of Noble Gases: Mendeleev’s periodic table did not include the noble gases, which were not yet discovered at the time. However, the table was later modified to include these elements in their own group.

Mendeleev’s periodic table was a significant milestone in the development of chemistry, and it provided a framework for understanding the properties of the elements. The periodic table is still used today as a fundamental tool for organizing the elements and predicting their properties.

Mendeleev’s Periodic Table was a groundbreaking achievement in the field of chemistry. Some of its major achievements are:

1. Prediction of New Elements: Mendeleev’s Periodic Table allowed him to predict the existence of several new elements, such as gallium, scandium, and germanium. He predicted their atomic weights, valences, and even some of their physical and chemical properties.
2. Filling of Gaps: The Periodic Table left gaps for elements that had not yet been discovered. When these elements were discovered, their properties closely matched Mendeleev’s predictions.
3. Classification of Elements: Mendeleev’s Periodic Table classified the elements into groups based on their properties, which made it easier to study and understand them. The Table also showed a clear trend in the properties of the elements across the rows and columns.
4. Correlation of Properties: The Periodic Table provided a way to correlate the properties of elements with their atomic structure, which helped to explain the behavior of elements in chemical reactions.
5. Development of Chemical Industry: The Periodic Table played an important role in the development of the chemical industry. It enabled scientists to identify elements that were useful in industrial processes and helped to develop new materials and compounds.
6. Confirmation of Atomic Theory: The Periodic Table provided evidence for the atomic theory of matter, which states that all matter is made up of atoms. The arrangement of the elements in the Table showed that the properties of the elements were determined by their atomic structure.

Mendeleev’s Periodic Table was a major milestone in the development of chemistry, and it provided a foundation for future research and discoveries in the field. It remains an essential tool for scientists today.

Mendeleev’s periodic table was a groundbreaking achievement in the field of chemistry. However, there were some limitations to his classification system. Some of the main limitations are:

1. Incomplete Arrangement of Elements: Mendeleev arranged elements based on their atomic weights, but some elements did not fit neatly into his periodic table. For example, argon, krypton, and xenon did not fit into his classification system because they have the same valence electron configuration as the noble gas in the previous period.
2. Anomalous Position of some elements: Mendeleev’s Periodic Table placed some elements in anomalous positions based on their properties. For example, iodine was placed in the same group as chlorine and bromine, even though it has a much larger atomic weight and a different valence electron configuration.
3. Ignored Isotopes: Mendeleev’s Periodic Table did not account for isotopes, which are atoms of the same element with different numbers of neutrons. Since isotopes have the same number of protons, they have the same atomic number and should be in the same position in the periodic table. However, they have different atomic weights, which would have affected Mendeleev’s placement of elements.
4. Inaccuracy in Atomic Weights: Mendeleev relied on the atomic weights of elements as a primary criterion for classification. However, the atomic weights available at the time were not always accurate. Some elements, such as cobalt and nickel, have nearly identical atomic weights, making it difficult to distinguish between them.
5. Ignored Rare Earth Elements: Mendeleev’s Periodic Table did not account for the rare earth elements, which are a group of elements that have similar properties. These elements were difficult to separate and study at the time, so Mendeleev did not include them in his classification system.

Despite these limitations, Mendeleev’s Periodic Table was a significant milestone in the development of chemistry and provided a foundation for future research and discoveries in the field.

Mendeleev’s periodic table was a groundbreaking

Two other elements that were discovered later and were predicted by Mendeleev’s periodic table are:

1. Scandium: Mendeleev left a gap between calcium and titanium, and he predicted that an element would be discovered with atomic weight around 44 that would fill the gap. In 1879, Lars Fredrik Nilson discovered scandium, which had an atomic weight of 45.
2. Germanium: Mendeleev also predicted that an element would be discovered between silicon and tin that would have an atomic weight of around 72. In 1886, Clemens Winkler discovered germanium, which had an atomic weight of 72.6.

Mendeleev used the following criteria in creating his periodic table:

1. Atomic Weight: Mendeleev arranged the elements in order of increasing atomic weight. He noticed that the properties of elements repeated periodically when they were arranged in this way.
2. Valence: Mendeleev also took into account the valence or combining power of the elements. He arranged elements with similar valences in the same group or column.
3. Chemical Properties: Mendeleev observed the chemical properties of the elements and grouped them together based on their similarities.
4. Physical Properties: Mendeleev also considered the physical properties of the elements, such as their melting points, boiling points, and densities, when grouping them.
5. Reactivity: Mendeleev also looked at the reactivity of elements and placed the most reactive elements on the left-hand side of the periodic table.
6. Spectral Lines: Mendeleev also considered the spectral lines of elements. He noticed that elements with similar properties had similar spectral lines.

Mendeleev used a combination of chemical and physical properties to arrange the elements in his periodic table, with a focus on atomic weight, valence, and periodicity of properties. His classification system laid the foundation for future research and development in the field of chemistry.

The noble gases are placed in a separate group because they have very stable electron configurations due to their filled outer electron shells. They have the maximum number of valence electrons possible for their respective energy levels, which makes them very unreactive and non-metallic. The noble gases do not readily form chemical bonds with other elements, making them inert or noble.

The unique properties of noble gases make them distinct from other elements, which is why they are placed in a separate group on the periodic table. In addition, they do not fit neatly into the traditional classification system based on valence electrons, as they have a full valence shell and do not form chemical bonds in the same way as other elements.

Placing the noble gases in a separate group also highlights the periodicity of their properties, as the group represents a complete set of elements with similar electronic configurations and properties. Overall, the placement of the noble gases in a separate group reflects their unique characteristics and allows for easy identification and classification of these elements.

The modern periodic table is an updated version of the periodic table created by Mendeleev. It arranges the elements in order of increasing atomic number, which is the number of protons in an atom’s nucleus. This arrangement reflects the periodic nature of the properties of elements and allows for easy identification and classification of elements.

The modern periodic table consists of rows, known as periods, and columns, known as groups. The rows represent the number of electron shells in an atom, while the columns represent elements with similar valence electron configurations and properties.

The modern periodic table has several advantages over Mendeleev’s periodic table, including:

1. It reflects the underlying structure of the atom: The modern periodic table arranges elements based on their atomic structure, which is fundamental to understanding the behavior of elements.
2. It includes all known elements: The modern periodic table includes all of the elements that have been discovered to date, whereas Mendeleev’s periodic table only included the elements known at the time.
3. It is based on empirical data: The modern periodic table is based on empirical data obtained through experimentation, whereas Mendeleev’s periodic table was based on observed patterns.
4. It allows for easy prediction of properties: The modern periodic table allows for easy prediction of the properties of elements based on their position in the table.

The modern periodic table is a fundamental tool in the field of chemistry, providing a systematic way to organize and understand the properties of elements.

The position of elements in the modern periodic table is based on their atomic number, which is the number of protons in the nucleus of an atom. The elements are arranged in rows, known as periods, and columns, known as groups, according to their electron configurations and properties.

Periods: There are seven periods in the modern periodic table, each representing the number of electron shells in an atom. The first period has only two elements, hydrogen and helium, while the remaining periods have up to 32 elements.

Groups: There are 18 groups in the modern periodic table, each representing elements with similar valence electron configurations and properties. The groups are numbered from 1 to 18, with the elements in Group 1 having one valence electron and the elements in Group 18 having a full valence shell.

The elements in the modern periodic table can be broadly classified into three categories based on their properties:

1. Metals: The majority of elements in the periodic table are metals. Metals are typically good conductors of heat and electricity, malleable, ductile, and have a metallic luster.
2. Nonmetals: Nonmetals are located on the right side of the periodic table and typically have poor conductivity, low melting and boiling points, and are not malleable or ductile. Some nonmetals, such as sulfur and iodine, are solids at room temperature, while others, such as chlorine and fluorine, are gases.
3. Metalloids: Metalloids, also known as semimetals, have properties intermediate between metals and nonmetals. They are located along the diagonal line between the metals and nonmetals in the periodic table.

The position of elements in the modern periodic table reflects their electronic structure and chemical properties, allowing for easy identification and classification of elements based on their position in the table.

The electronic configurations of the elements in the second period of the modern periodic table are as follows:

1. Lithium (Li): 1s^2 2s^1
2. Beryllium (Be): 1s^2 2s^2
3. Boron (B): 1s^2 2s^2 2p^1
4. Carbon (C): 1s^2 2s^2 2p^2
5. Nitrogen (N): 1s^2 2s^2 2p^3
6. Oxygen (O): 1s^2 2s^2 2p^4
7. Fluorine (F): 1s^2 2s^2 2p^5
8. Neon (Ne): 1s^2 2s^2 2p^6

These electronic configurations indicate the arrangement of electrons in the atom’s energy levels and sub-levels. The number of electrons in the outermost energy level, known as the valence shell, determines the chemical properties of an element. In the case of these elements in the second period, they all have two electrons in the first energy level and the remaining electrons in the second energy level. This makes them all similar in terms of their chemical behavior and properties.

No, these elements do not have the same number of valence electrons. The valence electrons are the electrons present in the outermost energy level of an atom, and they determine the chemical properties and reactivity of an element.

In the case of the elements in the second period of the modern periodic table (Li, Be, B, C, N, O, F, and Ne), they all have electrons in the second energy level, but the number of valence electrons varies from 1 to 8:

• Lithium (Li) has 1 valence electron.
• Beryllium (Be) has 2 valence electrons.
• Boron (B) has 3 valence electrons.
• Carbon (C) has 4 valence electrons.
• Nitrogen (N) has 5 valence electrons.
• Oxygen (O) has 6 valence electrons.
• Fluorine (F) has 7 valence electrons.
• Neon (Ne) has 8 valence electrons.

Therefore, even though these elements are located in the same period and have similar electronic configurations, they have different numbers of valence electrons and exhibit different chemical properties and reactivity.

No, the elements in the second period of the periodic table (Li, Be, B, C, N, O, F, and Ne) do not have the same number of shells.

The number of shells (energy levels) in an atom is determined by the principal quantum number (n) of the outermost electron. In this case, the outermost electron of these elements is located in the second energy level (n=2), but the number of electrons in the first energy level (n=1) varies from 1 to 2.

• Lithium (Li) has 2 energy levels, with 1 electron in the first energy level and 1 electron in the second energy level.

There are several trends that can be observed in the modern periodic table, including:

1. Atomic radius: The atomic radius decreases as you move from left to right across a period, and increases as you move down a group. This is because as you move across a period, the number of protons and electrons increases, which causes the attractive force between the positively charged nucleus and negatively charged electrons to increase, resulting in a smaller atomic radius. As you move down a group, the number of energy levels (shells) increases, which causes the atomic radius to increase.
2. Electronegativity: The electronegativity of an element increases as you move from left to right across a period, and decreases as you move down a group. This is because as you move across a period, the number of protons and electrons increases, which results in a greater attraction between the nucleus and electrons, making it harder for electrons to be removed from the atom. As you move down a group, the distance between the nucleus and outermost electrons increases, making it easier for electrons to be removed from the atom.
3. Ionization energy: The ionization energy of an element increases as you move from left to right across a period, and decreases as you move down a group. This is because as you move across a period, the number of protons and electrons increases, which results in a greater attraction between the nucleus and electrons, making it harder for electrons to be removed from the atom. As you move down a group, the distance between the nucleus and outermost electrons increases, making it easier for electrons to be removed from the atom.
4. Metallic character: The metallic character of an element decreases as you move from left to right across a period, and increases as you move down a group. This is because as you move across a period, the number of protons and electrons increases, which results in a greater attraction between the nucleus and electrons, making it harder for electrons to be removed from the atom. As you move down a group, the distance between the nucleus and outermost electrons increases, making it easier for electrons to be removed from the atom, resulting in increased metallic character.

These trends can be helpful in predicting the properties and behavior of elements in the periodic table.

Valency refers to the combining power of an element or its ability to form chemical bonds with other elements. It is determined by the number of electrons in the outermost shell (valence shell) of an atom.

The valence electrons of an atom are involved in chemical bonding and can be shared, gained, or lost during chemical reactions to form chemical bonds. The valence electrons are also responsible for determining the chemical and physical properties of an element, such as its reactivity, melting point, and boiling point.

The valency of an element can be determined by the number of valence electrons it has. For example, elements in group 1 of the periodic table (e.g. sodium, lithium) have 1 valence electron and a valency of 1, meaning they can form one chemical bond with another element. Elements in group 2 (e.g. magnesium, calcium) have 2 valence electrons and a valency of 2, meaning they can form two chemical bonds with other elements.

The valency of an element can also be determined by subtracting the number of valence electrons from 8 (for elements in the second period) or 18 (for elements in the third period or higher), as the valence shell can hold up to 8 electrons. For example, oxygen has 6 valence electrons, so its valency is 2 (8-6=2). Chlorine has 7 valence electrons, so its valency is 1 (8-7=1).

To calculate the valency of an element from its electronic configuration, you need to determine the number of valence electrons. Valence electrons are the electrons present in the outermost shell of an atom.

Here is a step-by-step method to determine the valency of an element from its electronic configuration:

1. Write down the electronic configuration of the element.
2. Determine the valence shell, which is the outermost shell that contains electrons.
3. Count the number of electrons in the valence shell. The number of electrons in the valence shell is equal to the group number of the element in the periodic table for main group elements.
4. Subtract the number of valence electrons from 8 (for elements in the second period) or 18 (for elements in the third period or higher) to get the valency.

Alternatively, you can also use the short-cut method to calculate the valency of an element:

1. Determine the group number of the element in the periodic table.
2. The group number of the element represents the number of valence electrons.
3. The valency of the element is equal to the group number for main group elements.

For example, the electronic configuration of sodium (Na) is 1s² 2s² 2p⁶ 3s¹. The valence shell is the 3s shell, which contains 1 electron. Therefore, the valency of sodium is 1.

Another example is the electronic configuration of oxygen (O) which is 1s² 2s² 2p⁴. The valence shell is the 2p shell, which contains 4 electrons. Therefore, the valency of oxygen is 2 (8-6=2).

Magnesium has an atomic number of 12, which means it has 12 electrons arranged in the electronic configuration: 1s² 2s² 2p⁶ 3s². The valence shell of magnesium is the 3s shell, which contains 2 electrons. Therefore, the valency of magnesium is 2.

Sulphur has an atomic number of 16, which means it has 16 electrons arranged in the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁴. The valence shell of sulphur is the 3p shell, which contains 6 electrons. Therefore, the valency of sulphur is 2, 4, or 6 depending on the reaction conditions. In most cases, sulphur has a valency of 2.

Here are the valencies of the first twenty elements:

1. Hydrogen (H) – 1
2. Helium (He) – 0
3. Lithium (Li) – 1
4. Beryllium (Be) – 2
5. Boron (B) – 3
6. Carbon (C) – 4
7. Nitrogen (N) – 3 or 5
8. Oxygen (O) – 2
9. Fluorine (F) – 1
10. Neon (Ne) – 0
11. Sodium (Na) – 1
12. Magnesium (Mg) – 2
13. Aluminum (Al) – 3
14. Silicon (Si) – 4
15. Phosphorus (P) – 3 or 5
16. Sulfur (S) – 2, 4, or 6
17. Chlorine (Cl) – 1
18. Argon (Ar) – 0
19. Potassium (K) – 1
20. Calcium (Ca) – 2

Note that the valency of some elements can vary depending on the reaction conditions. Therefore, the valencies listed here are the most common valencies observed for each element.

In a period of the periodic table, as we move from left to right, the valency of elements changes. The valency generally increases from 1 to 4 and then decreases back to 0 as we move across the period from left to right.

This happens because the number of valence electrons in the outermost shell of the atoms increases by one from left to right. The valence electrons determine the chemical properties of an element, so as the number of valence electrons increases, the valency of the element also increases.

In the first two periods of the periodic table, the valency generally follows the pattern of increasing from 1 to 4 and then decreasing back to 0. However, in the later periods, the pattern becomes more complex due to the involvement of the d and f orbitals in the valence electrons. Nonetheless, the general trend of increasing valency from left to right in a period holds true.

In a group of the periodic table, the valency of elements tends to remain the same as we move down the group. This is because the number of valence electrons in the outermost shell remains the same for all the elements in a group.

For example, the elements of Group 1, also known as the alkali metals, all have one valence electron in their outermost shell, so they have a valency of +1. Similarly, the elements of Group 2, also known as the alkaline earth metals, have two valence electrons in their outermost shell, so they have a valency of +2.

As we move down a group, the number of shells or energy levels in the atoms increases, but the outermost shell or valence shell remains the same. Therefore, the valency of elements in a group remains constant.

However, there are some exceptions to this trend, especially for the transition metals and the elements in the f-block. These elements have more complex electronic configurations and can have multiple valencies depending on the reaction conditions.

In general, atomic size increases as we move down a group in the periodic table. This is because the number of energy levels in the atom increases as we move down the group, leading to a greater distance between the nucleus and the outermost shell of electrons. For example, the atomic radius of the elements in Group 1 (Li, Na, K, etc.) increases as we move down the group.

Conversely, atomic size generally decreases as we move across a period from left to right. This is because, as we move across a period, the number of protons and electrons in the atoms increases, but the number of energy levels remains the same. This results in a greater pull on the electrons by the nucleus, making the atomic radius smaller. For example, the atomic radius of the elements in Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar) decreases as we move from left to right.

Exceptions to these general trends occur due to factors such as variations in electron configuration, nuclear charge, and shielding effect.

Arranging the given Period II elements in decreasing order of their atomic radii:

O (66 pm) < C (77 pm) < N (74 pm) < B (88 pm) < Be (111 pm) < Li (152 pm)

So the order of decreasing atomic radii is O, C, N, B, Be, Li.

No, the order of the elements given (O, C, N, B, Be, Li) does not correspond to the order of elements in a period in the modern periodic table. In the modern periodic table, the order of elements in the second period is:

Li, Be, B, C, N, O, F, Ne

Therefore, the given elements are not arranged in the pattern of a period in the periodic table.

The element with the largest atom is francium (Fr), which is located at the bottom of the alkali metal group (Group 1) in the periodic table. The atomic radius of francium is about 260 picometers (pm).

The element with the smallest atom is helium (He), which is located in the top right corner of the periodic table. The atomic radius of helium is about 31 pm.

As we move from left to right in a period of the modern periodic table, the atomic radius generally decreases. This is because, as we move across the period, the number of protons and electrons increases in the nucleus, which results in a greater attractive force on the electrons by the nucleus. Therefore, the electrons are pulled closer to the nucleus, and the atomic radius decreases.

This trend is observed across all the periods in the periodic table. However, there are a few exceptions where the atomic radius of some elements is larger than the element immediately preceding it in the same period. This happens because of the electron configuration of the elements, which affects the screening effect and the effective nuclear charge experienced by the outermost electrons.

The atomic radii of Group 1 elements in increasing order are:

Li < Na < K < Rb < Cs

So the order of atomic radii is:

152 pm (Li) < 186 pm (Na) < 231 pm (K) < 244 pm (Rb) < 262 pm (Cs)

Thus, lithium (Li) has the smallest atom, while cesium (Cs) has the largest atom among these Group 1 elements.

As we move down a group in the periodic table, the atomic radius generally increases. This is due to the increase in the number of shells occupied by the electrons as we move down the group. Each subsequent element in the group has an additional electron shell as compared to the previous element, which results in an increase in atomic size.

The increase in atomic size can be attributed to the shielding effect of inner electron shells, which reduces the effective nuclear charge experienced by the outermost electrons. This results in a weaker attraction between the nucleus and the outermost electrons, leading to an increase in atomic size.

The trend of increasing atomic size down the group is observed consistently in all the groups of the periodic table. However, there may be some anomalies in some groups where the trend is not strictly followed, which is due to the change in the screening effect or the effective nuclear charge.

Metallic Properties:

1. Lustrous: Metals have a characteristic shine or lustre which is due to their ability to reflect light.
2. Malleable: Metals can be beaten into thin sheets or flattened without breaking. This property is known as malleability.
3. Ductile: Metals can be drawn into wires. This property is known as ductility.
4. Good conductors of heat and electricity: Metals are good conductors of heat and electricity.
5. High melting and boiling points: Most metals have high melting and boiling points.

Non-metallic Properties:

1. Dull: Non-metals lack lustre and have a dull appearance.
2. Brittle: Non-metals are usually brittle and can break easily when subjected to force.
3. Insulators: Non-metals are poor conductors of heat and electricity.
4. Low melting and boiling points: Most non-metals have low melting and boiling points.
5. Non-malleable and non-ductile: Non-metals cannot be beaten into thin sheets or drawn into wires.

The metallic and non-metallic properties are generally found to be opposite to each other. While metals exhibit metallic properties, non-metals exhibit non-metallic properties. However, some elements may exhibit both metallic and non-metallic properties, such as metalloids, which have properties of both metals and non-metals.

1. Sodium (Na) – metal
2. Magnesium (Mg) – metal
3. Aluminum (Al) – metal
4. Silicon (Si) – non-metal
5. Phosphorus (P) – non-metal
6. Sulfur (S) – non-metal
7. Chlorine (Cl) – non-metal
8. Argon (Ar) – noble gas (non-metal)

Thus, the third period has 3 metals, 4 non-metals, and 1 noble gas.

Metals are found on the left side of the Periodic Table. They occupy the majority of the periodic table and are characterized by their shiny appearance, ability to conduct heat and electricity, malleability, ductility, and high tensile strength.

Non-metals are found on the right side of the Periodic Table. They include elements like hydrogen, helium, carbon, nitrogen, oxygen, fluorine, neon, chlorine, and argon. Non-metals are characterized by their dull appearance, poor ability to conduct heat and electricity, low melting and boiling points, and low densities.

The tendency to lose electrons generally decreases in a period as the atomic radius decreases and the effective nuclear charge increases, making it more difficult for the outermost electrons to be removed. This is because the increased nuclear charge attracts the electrons more strongly, making it harder for them to be removed. Therefore, elements towards the right side of a period tend to have a higher ionization energy and are less likely to lose electrons compared to elements towards the left side of the same period.

As we move from left to right across a period in the periodic table, the tendency of atoms to gain electrons increases. This is because the number of valence electrons in the outermost shell of the atoms increases, while the number of shells remains the same. As a result, the effective nuclear charge (i.e., the positive charge felt by the valence electrons) increases, making it more difficult for the atom to lose electrons and easier for it to gain electrons. Therefore, elements on the right side of a period are more likely to gain electrons than elements on the left side.

1. Based on atomic number: The Modern Periodic Table is based on the atomic number of elements, whereas Mendeléev’s table was based on atomic mass. By arranging the elements in the increasing order of atomic number, the Modern Periodic Table was able to overcome the inconsistencies that arose due to the anomalous position of certain elements based on their atomic mass.
2. Position of isotopes: The Modern Periodic Table also takes into account the presence of isotopes of an element. Isotopes have the same atomic number but different atomic masses. This was not considered in Mendeléev’s table, which created inconsistencies in the position of elements.
3. Placement of noble gases: Mendeléev’s Periodic Table did not include noble gases as they were not discovered yet. However, when noble gases were discovered, they did not fit into any of Mendeléev’s groups. In the Modern Periodic Table, a separate group has been created for noble gases, which removes the anomaly.
4. Inclusion of lanthanides and actinides: Mendeléev’s Periodic Table did not have a proper place for the lanthanides and actinides. However, in the Modern Periodic Table, these elements have been placed separately at the bottom of the main table, which removes the anomaly.

Overall, the Modern Periodic Table has been able to provide a more systematic and accurate representation of the elements, and has eliminated many of the inconsistencies and anomalies that were present in Mendeléev’s table.

Two elements that are expected to show chemical reactions similar to magnesium are calcium (Ca) and strontium (Sr).

This is because calcium and strontium are in the same group as magnesium, i.e., group 2, which is also known as the alkaline earth metals. Elements in the same group have similar chemical properties due to the same number of valence electrons in their outermost shell. Magnesium, calcium, and strontium all have two valence electrons, which makes them likely to form similar types of chemical bonds and compounds.

(a) Lithium (Li), Sodium (Na), Potassium (K) (b) Calcium (Ca), Strontium (Sr) (c) Helium (He), Neon (Ne), Argon (Ar)

Yes, there is a similarity in the atoms of lithium, sodium, and potassium as they all belong to Group 1 of the periodic table, also known as the alkali metals. They all have a single electron in their outermost shell, which makes them highly reactive and prone to losing that electron to form a positively charged ion. This similarity in electronic configuration is what makes them react similarly with water to liberate hydrogen gas.

Among the first ten elements in the Modern Periodic Table, the metals are:

1. Lithium (Li)
2. Beryllium (Be)
3. Sodium (Na)
4. Magnesium (Mg)
5. Aluminum (Al)

Beryllium (Be) would be expected to have maximum metallic character among the given elements. This is because Be is located in Group 2 of the periodic table, which is known as the Alkaline Earth Metals group. The elements in this group are highly metallic in nature, with low ionization energies and low electronegativities. In contrast, the other elements listed (Ga, Ge, As, and Se) are located in Groups 13, 14, and 15, which are less metallic and have higher ionization energies and electronegativities.

(b) The number of valence electrons increases is not a correct statement about the trends when going from left to right across the periods of periodic Table. The number of valence electrons remains the same for elements belonging to the same period.

The correct answer is (b) Mg. The element X is most likely a metal since it forms a chloride with a high melting point. Also, the formula of the chloride is XCl2, which indicates that X has a valency of 2. Magnesium (Mg) is the only option in the given list that has a valency of 2 and is a metal.

(a) The element with two shells, both of which are completely filled with electrons is Helium (He). (b) The element with the electronic configuration 2, 8, 2 is Calcium (Ca). (c) The element with a total of three shells, with four electrons in its valence shell is Silicon (Si). (d) The element with a total of two shells, with three electrons in its valence shell is Boron (B). (e) The element with twice as many electrons in its second shell as in its first shell is Carbon (C).

(a) All elements in the same column of the Periodic Table as boron have the same number of valence electrons, which is three. (b) All elements in the same column of the Periodic Table as fluorine have the same number of valence electrons, which is seven. Additionally, they tend to form negatively charged ions (anions) in chemical reactions.

(a) The atomic number of the element is 17.

(b) The element with the electronic configuration of 2, 8, 7 is chlorine (Cl), which is in the same group as fluorine (F) with an atomic number of 9. Therefore, chlorine would be chemically similar to fluorine.

The electronic configuration of nitrogen is 1s2 2s2 2p3 and that of phosphorus is 1s2 2s2 2p6 3s2 3p3.

Nitrogen has 5 valence electrons and phosphorus has 5 valence electrons as well. Both these elements require three electrons to complete their octet configuration. However, phosphorus has a larger atomic radius than nitrogen, which means that the valence electrons are farther from the nucleus and thus, it is easier for phosphorus to attract an electron towards itself. Therefore, phosphorus is more electronegative than nitrogen.

The element with atomic number 12 (magnesium) and the element with atomic number 38 (strontium) are the closest to calcium in terms of physical and chemical properties, as they all belong to the same group (Group 2) of the Modern Periodic Table. These elements have similar electronic configurations and exhibit similar chemical reactions due to their tendency to lose two electrons and form divalent cations. Additionally, they are all metals with similar physical properties such as high melting and boiling points and good conductivity. The element with atomic number 19 (potassium) is in a different group and has a different electronic configuration, while the element with atomic number 21 (scandium) is a transition metal with very different properties from those of calcium.

Mendeléev’s Periodic Table and the Modern Periodic Table are two different arrangements of the elements, but they share some similarities. Here are some of the main differences and similarities:

Mendeléev’s Periodic Table:

• It was developed by Dmitri Mendeléev in 1869.
• The elements were arranged in order of increasing atomic mass.
• The elements were grouped into columns (groups) and rows (periods) based on their chemical and physical properties.
• The gaps were left for undiscovered elements that Mendeléev predicted would be found.
• There were only 63 known elements at the time of its creation.

Modern Periodic Table:

• It was developed by Henry Moseley in 1913.
• The elements are arranged in order of increasing atomic number.
• The elements are grouped into columns (groups) and rows (periods) based on their chemical and physical properties.
• There are no gaps for undiscovered elements because all known elements are accounted for.
• There are currently 118 known elements.

Some similarities between the two tables include:

• Both tables group elements based on their properties.
• Both tables have columns (groups) and rows (periods).
• Both tables have a similar overall shape.

Some differences between the two tables include:

• The order of the elements is based on atomic mass in Mendeléev’s Periodic Table and atomic number in the Modern Periodic Table.
• The Modern Periodic Table has more known elements than Mendeléev’s Periodic Table.
• The Modern Periodic Table is more precise and accurate due to advances in technology and our understanding of atomic structure.

• Chapter 5: Periodic Classification of Elements