Structure of the periodic table
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Intros
Examples
Lessons
- Apply knowledge of the periodic table structure to classify elements.
For each chemical element, state the chemical group it is in, whether it is a metal or non-metal, and the block of the periodic table in which it belongs. - Apply your knowledge of the structure of the periodic table to identify elements of similar properties.
For each chemical element, state another chemical element which shares similar properties to it.
Topic Notes
Introduction to the Structure of the Periodic Table
Welcome to our exploration of the periodic table's structure! This organized layout of elements is a cornerstone of chemistry, providing a wealth of information at a glance. The periodic table arranges elements based on their atomic number, electron configuration, and recurring chemical properties. It's divided into periods (rows) and groups (columns), each with specific characteristics. As we delve into this topic, you'll discover how this structure helps predict element behavior and relationships. To kick things off, we've prepared an introduction video that will give you a visual overview of the periodic table. This video is a great starting point to understand the basics before we dive deeper. It'll help you grasp the periodic table's importance in chemistry and how its structure reflects the fundamental properties of elements. Ready to unravel the mysteries of the periodic table? Let's begin this exciting journey together!
Historical Development of the Periodic Table
Early Challenges in Organizing Elements
The journey to create the periodic table we know today was far from straightforward. Initially, scientists faced numerous challenges in organizing the known elements in a logical and coherent manner. The primary reason why the periodic table initially did not follow a pattern was the limited understanding of atomic structure and properties.
In the early 19th century, chemists were aware of several dozen elements but struggled to find a systematic way to arrange them. They noticed similarities between certain elements but couldn't explain these relationships scientifically. Some early attempts at classification included Johann Döbereiner's "triads" and John Newlands' "Law of Octaves," but these systems were limited and didn't accommodate all known elements.
Mendeleev's Groundbreaking Contribution
The most significant breakthrough came with Dmitri Mendeleev's contribution to the periodic table in 1869. Mendeleev, a Russian chemist, approached the problem differently. He created cards for each of the 63 known elements, listing their properties and atomic weights. Through careful analysis, he noticed patterns emerging when elements were arranged by increasing atomic mass.
Mendeleev's genius lay in his ability to:
- Arrange elements in order of increasing atomic mass
- Group elements with similar properties together
- Leave gaps for undiscovered elements
- Predict properties of these unknown elements
This approach allowed Mendeleev to create a table that not only organized known elements but also predicted the existence and properties of elements yet to be discovered. His predictions, such as for gallium, germanium, and scandium, were later proven correct, cementing the validity of his periodic system.
Evolution of the Periodic Table
Despite Mendeleev's breakthrough, the early periodic table still had inconsistencies. Some elements seemed out of place when arranged by atomic mass. For instance, tellurium (atomic mass 127.6) was placed before iodine (atomic mass 126.9) to maintain chemical property patterns, contradicting the atomic mass order.
The resolution to these inconsistencies in the early periodic table came with the discovery of the atomic structure, particularly the concept of atomic number (number of protons in an atom's nucleus). In 1913, Henry Moseley demonstrated that the atomic number, not atomic mass, was the fundamental property determining an element's chemical behavior.
This discovery led to the modern periodic table, where elements are arranged by increasing atomic number. This arrangement resolved the previous anomalies and provided a more accurate representation of elemental relationships.
Why the Initial Table Didn't Follow a Perfect Pattern
Several factors contributed to the initial lack of a clear pattern in the periodic table:
- Incomplete knowledge of elements: Many elements were yet to be discovered, leaving gaps in the pattern.
- Reliance on atomic mass: Using atomic mass as the primary organizing principle led to some misplacements.
- Limited understanding of atomic structure: The concept of protons and atomic number was unknown.
- Isotopes: The existence of isotopes (atoms of the same element with different atomic masses) wasn't yet understood, causing confusion in mass-based arrangements.
Conclusion
The development of the periodic table is a testament to scientific progress and the power of systematic observation. From the initial challenges of organizing seemingly disparate elements to Mendeleev's revolutionary approach and the later refinements based on atomic structure, the periodic table has evolved into an indispensable tool in chemistry. Today, it stands as a perfect example of how scientific understanding grows and adapts with new discoveries, providing a clear, logical arrangement of elements based on their fundamental properties.
Basic Structure of the Modern Periodic Table
Overview of the Periodic Table
The modern periodic table is a cornerstone of chemistry, providing a systematic arrangement of all known chemical elements. This ingenious layout offers a wealth of information about element properties and their relationships, making it an indispensable tool for scientists, students, and enthusiasts alike.
Groups: The Vertical Columns
One of the key features of the periodic table structure is its vertical columns, known as groups. There are 18 groups in total, each designated by a number from 1 to 18. Elements within the same group share similar chemical properties due to their identical outer electron configurations. For instance, Group 1 elements, also called alkali metals, are highly reactive and form strong bases when combined with water.
Periods: The Horizontal Rows
The horizontal rows of the periodic table are called periods. There are seven periods in total, with each new period representing a new electron shell in the atoms of those elements. As you move from left to right across a period, the atomic number increases by one, and elements generally become less metallic and more nonmetallic in nature.
Blocks: Regions of the Periodic Table
The periodic table is divided into four distinct blocks: s-block, p-block, d-block, and f-block. These periodic table blocks correspond to the type of subshell that is being filled with electrons in the elements' atoms. This arrangement helps in understanding the electronic configuration and chemical behavior of elements.
Element Properties and Periodic Trends
The periodic table's structure allows for the observation of trends in element properties. As you move down a group, atomic size trends generally increase due to the addition of new electron shells. Across a period, atomic size trends typically decrease as the nuclear charge increases, pulling electrons closer to the nucleus. Other properties that follow periodic trends include ionization energy trends, electron affinity, and electronegativity.
Metals, Nonmetals, and Metalloids
The periodic table clearly delineates elements into metals, nonmetals, and metalloids. Metals, found on the left and center of the table, are generally good conductors of heat and electricity. Nonmetals, located on the right side, typically have opposite properties to metals. Metalloids, straddling the boundary between metals and nonmetals, exhibit properties of both.
Significance of the Periodic Table Arrangement
The arrangement of elements in the periodic table is not arbitrary; it reflects fundamental principles of atomic structure and quantum mechanics. This organization allows scientists to predict the properties of elements, even those that have not yet been discovered. The periodic table's layout has proven invaluable in understanding chemical bonding, reactivity, and the behavior of elements in various compounds.
Modern Applications and Ongoing Research
While the basic structure of the periodic table has remained largely unchanged since Mendeleev's time, ongoing research continues to refine our understanding of element properties and relationships. The discovery of new synthetic elements and the study of exotic isotopes push the boundaries of the table, leading to potential expansions and modifications in the future.
Conclusion
The modern periodic table's structure, with its groups, periods, and blocks, provides a comprehensive framework for understanding the chemical elements. Its arrangement not only organizes elements based on their properties but also reveals fundamental truths about atomic structure and chemical behavior. As a vital tool in chemistry and related sciences, the periodic table continues to guide research, education, and innovation in countless fields, from materials science to nuclear physics.
Blocks in the Periodic Table
The periodic table is organized into four distinct blocks: s-block, p-block, d-block, and f-block. These blocks are named after the subshell in which the last electron of an element is placed according to its electron configuration. Understanding these blocks is crucial for comprehending element properties and chemical behavior.
The s-block consists of the first two groups (columns) of the periodic table. It includes alkali metals (Group 1) and alkaline earth metals (Group 2), as well as hydrogen and helium. Elements in this block have their outermost electrons in the s subshell. For example, sodium (Na) has the electron configuration [Ne]3s1, with its last electron in the 3s orbital. S-block elements are highly reactive, with alkali metals being the most reactive of all elements.
The p-block encompasses groups 13 to 18 and includes familiar elements like carbon, nitrogen, and oxygen. These elements have their outermost electrons in the p subshell. For instance, chlorine (Cl) has the configuration [Ne]3s23p5. P-block elements display a wide range of properties, from metals to nonmetals and noble gases. They often form covalent bonds and can exhibit multiple oxidation states.
The d-block, also known as the transition metals, comprises groups 3 to 12. These elements have partially filled d orbitals in their electron configurations. For example, iron (Fe) has the configuration [Ar]3d64s2. D-block elements are characterized by their ability to form colored compounds, exhibit multiple oxidation states, and act as effective catalysts. They also tend to form complex ions and have high melting points.
The f-block consists of two series: lanthanides and actinides, often placed separately at the bottom of the periodic table. These elements have electrons filling the f subshell. For instance, cerium (Ce) has the configuration [Xe]4f15d16s2. F-block elements are known for their unique magnetic and spectroscopic properties. They are less reactive than s-block elements but more reactive than d-block elements.
The electron configuration of an element directly relates to its position in these blocks. As we move across a period (row) in the periodic table, electrons fill orbitals in the order: s, p, d, and then f. This filling pattern, known as the Aufbau principle, explains the structure of the periodic table and the characteristics of each block.
The s-block elements, with their simple configurations, tend to lose electrons easily, making them highly reactive. P-block elements, with more complex outer shell structures, exhibit a range of bonding behaviors. D-block elements, with their partially filled d orbitals, show variable oxidation states and form colorful compounds. F-block elements, with their unique f orbitals, display distinctive spectroscopic and magnetic properties.
These block characteristics significantly influence element properties. For example, the reactivity of s-block elements decreases down the group due to increasing atomic size and decreasing ionization energy. In the p-block, we observe trends in electronegativity and atomic radius across periods and down groups. D-block elements show less pronounced trends due to the influence of d orbitals, leading to their diverse chemical behaviors.
Understanding these blocks helps predict and explain various chemical and physical properties. For instance, the high melting points of d-block elements can be attributed to strong metallic bonding involving d electrons. The color of transition metal compounds is often due to d-d electron transitions. The magnetic properties of f-block elements are a result of unpaired f electrons.
In conclusion, the s, p, d, and f blocks of the periodic table provide a fundamental framework for understanding element properties based on electron configuration. This organization not only helps in predicting chemical behavior but also in explaining trends across the periodic table. Whether it's the high reactivity of alkali metals in the s-block, the diverse bonding in the p-block, the colorful compounds of the d-block, or the unique properties of f-block elements, each block contributes to the rich tapestry of chemical elements that make up our world.
Metals, Non-metals, and Metalloids
The periodic table organization is a treasure trove of information about chemical elements, neatly organizing them based on their properties. One of the most fundamental ways to categorize elements is by dividing them into metals, non-metals, and metalloids. This classification provides valuable insights into the behavior and characteristics of elements, making it essential for students and professionals alike to understand their distribution and properties.
Let's start with metals, which dominate the periodic table. These elements are typically found on the left side and in the middle of the table. Some well-known examples include iron (Fe), gold (Au), and sodium (Na). Metals are characterized by their lustrous appearance, high electrical and thermal conductivity, and malleability. They tend to lose electrons easily, forming positive ions, which explains their reactivity with non-metals.
Moving to the right side of the periodic table, we encounter the non-metals. These elements, such as oxygen (O), chlorine (Cl), and carbon (C), exhibit properties quite different from metals. Non-metals are generally poor conductors of heat and electricity, and they tend to gain electrons in chemical reactions, forming negative ions. Unlike the shiny metals, non-metals can be dull in appearance and are often found in gaseous or brittle solid states at room temperature.
Bridging the gap between metals and non-metals are the fascinating metalloids. These elements, including silicon (Si), boron (B), and germanium (Ge), are found along a zigzag line that separates metals from non-metals on the periodic table. Metalloids possess properties of both metals and non-metals, making them unique and valuable in various applications, particularly in the semiconductor industry.
The distribution of these element types across the periodic table organization follows a pattern. As we move from left to right across a period (row), elements generally transition from metals to metalloids to non-metals. For instance, in period 3, we start with the reactive metal sodium, progress through the metalloid silicon, and end with the non-metal chlorine. This trend is due to the increasing number of electrons in the outermost shell, which affects how easily elements can lose or gain electrons.
Vertically, elements in the same group (column) tend to have similar properties. The leftmost groups are home to the most reactive metals, such as the alkali metals reactivity (Group 1) and alkaline earth metals (Group 2). On the far right, we find the noble gases stability (Group 18), which are non-metals known for their stability and low reactivity.
Understanding the properties of metals, non-metals, and metalloids is crucial in various fields. In materials science, the conductivity of metals is harnessed for electrical wiring, while the semiconducting properties of metalloids like silicon form the basis of modern electronics. Non-metals play vital roles in biological processes; for example, oxygen is essential for respiration, and carbon forms the backbone of organic compounds.
The periodic table's organization allows us to predict element properties based on their position. For instance, elements in the same group often form similar compounds. The halogens (Group 17), including fluorine (F) and chlorine (Cl), readily form salts with metals. Meanwhile, transition metals, located in the middle of the table, often exhibit multiple oxidation states, making them valuable in catalysis and complex ion formation.
As we delve deeper into the periodic table, we discover fascinating exceptions and nuances. For example, hydrogen (H), despite being placed with the alkali metals, behaves more like a non-metal in many situations. Some elements, like polonium (Po), display properties that blur the lines between categories, showcasing the complexity and richness of elemental chemistry.
In conclusion, the distribution of metals, non-metals, and metalloids in the periodic table provides a fundamental framework for understanding element properties. This organization not only helps in predicting chemical behavior but also guides research and applications across various scientific and technological fields. By appreciating these distinctions, we gain a deeper insight into the building blocks of our universe and their diverse roles in shaping the world around us.
Trends in the Periodic Table
The periodic table trends is a treasure trove of information, revealing fascinating trends in elemental properties. These periodic table trends provide crucial insights into atomic structure and chemical behavior. In this section, we'll explore three major trends: atomic radius, ionization energy, and electronegativity trends.
Atomic Radius
Atomic radius refers to the size of an atom, typically measured from the nucleus to the outermost electron shell. As we move from left to right across a period in the periodic table, the atomic radius generally decreases. This trend occurs because each successive element has one more proton in its nucleus, increasing the nuclear charge. The stronger nuclear attraction pulls the electrons closer to the nucleus, resulting in a smaller atom. Conversely, as we move down a group, the atomic radius increases. This is due to the addition of new electron shells, which are farther from the nucleus.
Ionization Energy
Ionization energy is the amount of energy required to remove an electron from an atom in its gaseous state. This property exhibits a trend opposite to atomic radius. As we move from left to right across a period, ionization energy generally increases. This is because the electrons are held more tightly by the increased nuclear charge, requiring more energy to remove them. Moving down a group, ionization energy typically decreases. The outermost electrons in larger atoms are farther from the nucleus and less tightly bound, making them easier to remove.
Electronegativity
Electronegativity trends measures an atom's ability to attract electrons in a chemical bond. This property follows a similar trend to ionization energy. Electronegativity trends increases from left to right across a period due to the increasing nuclear charge and decreasing atomic size. Elements on the right side of the periodic table, such as fluorine and chlorine, have the highest electronegativity values. As we move down a group, electronegativity generally decreases because the valence electrons are farther from the nucleus and less strongly attracted.
Illustrating the Trends
To visualize these trends, imagine arrows on the periodic table: - Atomic radius: Decreasing arrow pointing right across periods, increasing arrow pointing down groups. - Ionization energy: Increasing arrow pointing right across periods, decreasing arrow pointing down groups. - Electronegativity: Increasing arrow pointing right across periods, decreasing arrow pointing down groups.
Examples and Exceptions
While these trends are generally consistent, there are some exceptions: - Atomic radius: The transition metals may not follow the trend strictly due to their complex electron configurations. - Ionization energy: Elements like beryllium (Be) and nitrogen (N) have unexpectedly high ionization energy due to their electron configurations. - Electronegativity: Noble gases are typically excluded from electronegativity trends scales due to their low reactivity.
Practical Applications
Understanding these periodic trends is crucial for predicting chemical behavior and reactivity. For instance: - Atomic radius influences the strength of chemical bonds and the size of molecules. - Ionization energy helps predict the ease of forming cations and the strength of metallic character. - Electronegativity trends is essential for understanding bond polarity and the nature of chemical reactions.
Conclusion
The periodic table's trends in atomic radius, ionization energy, and electronegativity trends provide a powerful framework for understanding elemental properties. By recognizing these patterns, chemists can make predictions about chemical behavior, design new materials, and unravel the mysteries of atomic interactions. As you explore the periodic table, remember that these trends are not just abstract concepts but fundamental principles that shape the world of chemistry and materials science.
Conclusion
Understanding the structure of the periodic table is crucial for anyone studying chemistry. This organized layout of elements provides a wealth of information at a glance. The table's rows and columns reveal patterns in atomic structure, electron configuration, and chemical properties. Elements are arranged by increasing atomic number, with metals, nonmetals, and metalloids grouped together. The periodic table's structure allows scientists to predict element behavior and discover new elements. Its importance in chemistry cannot be overstated, as it forms the foundation for understanding chemical reactions and compound formation. We encourage you to explore this topic further, as mastering the periodic table will greatly enhance your chemistry knowledge. The introduction video provides an excellent starting point for your journey into the fascinating world of elements. Remember, the periodic table is not just a chart; it's a powerful tool that unlocks the secrets of matter itself. Keep exploring and let your curiosity guide you through the elements!
Using the Periodic Table: Changing Mendeleev's Table
In this guide, we will explore the structure of the periodic table and how it has evolved from Mendeleev's original design. We will discuss the key changes made to the table, the reasons behind these changes, and how the modern periodic table is organized.
Step 1: Understanding Mendeleev's Contribution
Dmitri Mendeleev made a significant contribution to the development of the periodic table. He organized the elements based on their atomic mass and grouped them into columns (groups or families) and rows (periods). Mendeleev's table highlighted the periodic nature of the elements, where properties repeated in a predictable pattern across different rows. He also left gaps in his table, predicting the existence of elements that had not yet been discovered, such as gallium and scandium.
Step 2: Identifying the Limitations of Atomic Mass
While Mendeleev's table was groundbreaking, it had some limitations. One major issue was the inconsistency in measuring atomic mass, which made it difficult for scientists to reach a consensus. Additionally, some elements did not fit well into the table when ordered strictly by atomic mass. For example, argon and potassium had masses that did not align with their chemical properties, leading to confusion and the need for a new ordering method.
Step 3: Introduction of Proton Number
As scientific understanding of the atom improved, particularly with the work of J.J. Thompson and Ernest Rutherford, it became clear that ordering elements by their proton number (atomic number) was more effective. The proton number represents the number of protons in an atom's nucleus and provides a more consistent basis for organizing the elements. This change resolved the inconsistencies observed with atomic mass and better reflected the periodic nature of the elements.
Step 4: The Modern Periodic Table
The modern periodic table is now ordered by proton number rather than atomic mass. This arrangement ensures that elements with similar chemical properties are grouped together in columns, while the rows reflect the repeating patterns of these properties. The periodic law, which states that the properties of elements repeat when the table is ordered by proton number, underpins the structure of the modern periodic table. This law highlights the periodic (repeating) nature of the elements, which is the fundamental principle behind the table's design.
Step 5: Conclusion
In summary, the periodic table has evolved significantly since Mendeleev's time. The shift from ordering by atomic mass to proton number has provided a more accurate and consistent framework for understanding the elements and their properties. The modern periodic table continues to be an essential tool in chemistry, reflecting the periodic law and the repeating nature of elemental properties.
FAQs
Q1: What was wrong with the original periodic table?
A1: The original periodic table, developed by Mendeleev, had some inconsistencies due to arranging elements by atomic mass rather than atomic number. This led to some elements being placed out of order based on their chemical properties. For example, tellurium was placed before iodine despite having a higher atomic mass. These issues were resolved when the table was reorganized based on atomic number.
Q2: What is the basic structure of a periodic table?
A2: The basic structure of the periodic table consists of rows (periods) and columns (groups). Elements are arranged in order of increasing atomic number from left to right and top to bottom. The table is divided into four blocks: s, p, d, and f, based on electron configurations. It also separates elements into metals, nonmetals, and metalloids.
Q3: How do you summarize the periodic table?
A3: The periodic table is a systematic arrangement of chemical elements, organized by their atomic number, electron configuration, and recurring chemical properties. It consists of 18 groups (columns) and 7 periods (rows), with elements in the same group sharing similar characteristics. The table provides essential information about each element, including its symbol, atomic number, and atomic mass.
Q4: What are the parts of the periodic table? Describe each of them.
A4: The main parts of the periodic table are:
1. Groups (columns): Elements with similar properties.
2. Periods (rows): Elements with the same number of electron shells.
3. Atomic number: The number of protons in an atom's nucleus.
4. Element symbol: A one or two-letter abbreviation for each element.
5. Atomic mass: The average mass of an atom of the element.
6. Blocks: s, p, d, and f, based on electron configurations.
Q5: What are the major trends in the periodic table?
A5: The major trends in the periodic table include:
1. Atomic radius: Decreases from left to right across a period and increases down a group.
2. Ionization energy: Increases from left to right across a period and decreases down a group.
3. Electronegativity: Increases from left to right across a period and decreases down a group.
4. Metallic character: Decreases from left to right across a period and increases down a group.
Prerequisite Topics
Understanding the structure of the periodic table is a fundamental aspect of chemistry, but to truly grasp its significance, it's crucial to have a solid foundation in several prerequisite topics. These topics provide the necessary context and background knowledge to fully appreciate the periodic table's organization and its implications for chemical behavior.
One of the most important prerequisites is understanding the properties of elements in the periodic table. This knowledge forms the basis for recognizing patterns and trends across the table. By familiarizing yourself with the chemical properties of elements, you'll be better equipped to predict how elements will interact and behave in various chemical reactions.
To appreciate the current structure of the periodic table, it's valuable to explore the history and development of the periodic table. This historical perspective helps you understand why the table is organized the way it is and how scientists discovered the periodic law. Learning about the periodic table periods provides insight into the table's rows and their significance.
A key aspect of the periodic table's structure is the organization of elements into groups. Studying specific groups, such as the Group 1 and Group 2 elements (alkali and alkaline earth metals), helps illustrate how elements with similar properties are grouped together. Understanding alkali metals' reactivity, for example, demonstrates how the table's structure reflects chemical behavior.
The periodic table's structure is closely tied to various trends in elemental properties. One crucial trend to understand is ionization energy. Grasping ionization energy trends across the table provides insight into how easily atoms lose electrons, which is fundamental to understanding chemical bonding and reactivity.
Similarly, electronegativity is another vital concept related to the periodic table's structure. Recognizing electronegativity trends helps predict the nature of chemical bonds between elements and explains many patterns in chemical behavior across the table.
Lastly, understanding atomic radius trends is essential for comprehending the periodic table's structure. The changes in atomic size across periods and down groups are directly related to the table's organization and influence many other elemental properties.
By mastering these prerequisite topics, you'll develop a comprehensive understanding of why the periodic table is structured as it is. This knowledge will enable you to navigate the table more effectively, predict elemental properties, and understand chemical behavior on a deeper level. The interconnectedness of these concepts with the periodic table's structure highlights the importance of building a strong foundation in these areas before delving into more advanced chemical studies.
- The changes made to the Periodic Table since Mendeleev's major contribution.
- The basic structure and layout of the modern Periodic Table.
- The precise grouping and arrangement of the Periodic Table.
Notes:
- As seen in the lesson on History and Development of the Periodic Table, one of the main reasons the development of the Periodic Table was initially slow is because data quality was often poor. Inaccurate measurements meant organizing elements correctly was hard.
- As improvements meant more and better data could be used, some changes were made to Mendeleev's Periodic Table. One big change was in ordering the table by atomic number (proton number) rather than by atomic mass. This solved the problem of isotopes; when arranged by atomic mass, some elements look like they are in the wrong place in the table (e.g. K has a lower atomic mass than Ar, even though K has one more proton). The current periodic table, then, obeys the Periodic Law: The properties of the elements repeat periodically when ordered according to their atomic numbers.
- The basic layout of the current periodic table has:
- Columns known as groups, the elements inside each of which having similar properties.
- Rows known as periods. The elements in a given period have the same number of electron shells.
- The current periodic table is arranged in a way that separates metals and non-metals:
- Metals are found on the left hand side of the periodic table.
- Non-metals are found on the right hand side of the periodic table.
- Metalloids or semiconductors are found on the border between metals and non-metals. These elements have a mix of metal and non-metal properties.
- About 80% of the elements in the periodic table are metals. The 'border' between metals and non-metals begins with boron and runs diagonally down to between non-metal Astatine (At) and Polonium metal (Po).
- The current periodic table has its specific shape to clearly show the different "blocks" of elements, in terms of their electron subshells (see lesson Electronic structure: Subshells).
- The first two groups (columns) form the s-block, the taller section on the left. This contains the alkali metals (group 1) and the alkali earth metals (group 2).
- The next ten groups form the d-block, generally known as the transition metals. This is the central block where the table 'dips'.
- The next six groups, where the table rises up again, is the p-block. It contains the metalloids and non-metals, including the halogens and the noble gases.
- The 'island' normally shown alone at the bottom of the periodic table is the f-block. It is actually an insert, inside the lower part of the d-block. It contains two rows known as the lanthanides and actinides.
- The periodic table has its specific shape to indicate the electron sub-shells and the block which a given element falls into. This is based on experimental data; the energy cost of removing the first outer shell electron (the ionization energy) from a sample of a given element follows a consistent pattern across the period (see lesson Periodic Trends: Ionization energy). This pattern led chemists to develop the sub-shell theory where not all electrons in a shell are in the same state – some electrons were in s orbitals, others p, d or f orbitals. The outer electrons occupying these orbitals, as suggested by experimental data, is represented in the shape of the periodic table we use.
