First 20 Elements of the Periodic Table: Names, Symbols & Atomic Numbers

Edited By Team Careers360 | Updated on Jul 02, 2025 04:34 PM IST

The first 20 elements of the Periodic table form the core of Chemistry. Every element in the periodic table is arranged based on its chemical and physical properties, such as atomic number, electronic configuration, and recurring chemical properties. These are the very basic elements, but they form the basis for the entire chemistry.
These elements represent the wide range of Metals, Non-metals, and Noble Gases. From the importance of Carbon in Organic life to the essential role of Calcium in bones and teeth, these elements play an important role in everyday life. Inside the periodic table , elements are arranged in the form of rows and columns. The rows are called periods, and the columns are called groups. The basic understanding of these elements helps to build a strong foundation in Chemistry.

Atomic Numbers and Symbols of the First 20 Elements

Atomic numbers	Element	Symbol
	Hydrogen	H
	Helium	He
	Lithium	Li
	Berilliyum	Be
	Boron	B
	Carbon	C
	Nitrogen	N
	Oxygen	O
	Fluorine	F
	Neon	Ne
	Sodium	Na
	Magnesium	Mg
	Aluminium	Al
	Silicon	Si
	Phosphorous	P
	Sulfur	S
	Chlorine	Cl
	Argon	Ar
	Potassium	K
	calcium	Ca

An in-depth look at the first 20 elements in the periodic table is useful. Six of these elements make up nearly all of the mass of the human body. The first 20 elements provide an excellent overview of the various element groups. They can also be found in more common chemical processes. The elements are given in ascending atomic numbers from 1 to 20 order. The atomic number refers to the number of protons in each element's atoms.

Also read -

NCERT Solutions for Class 11 Chemistry	NCERT notes Class 11 Chemistry Chapter 3 Classification of Elements and Periodicity in Properties
NCERT Solutions for Class 12 Chemistry	NCERT Solutions for Class 11 Chemistry Chapter 3 Classification of Elements and Periodicity in Properties
NCERT Solutions for All Subjects	NCERT Exemplar Class 11 Chemistry Solutions Chapter 3 Classification of Elements and Periodicity in Properties

First 20 elements Of The Periodic Table

The periodic table is the representation of elements in the increasing order of their atomic number.

The Periodic table

Name and Symbols of Elements

The atomic number, element name, or element symbol can all be used to identify elements. A one- or two-letter abbreviation of the name serves as the emblem. Some symbols, on the other hand, allude to old element names. The symbol for sodium, for given as Na. This is a reference to the Latin word natrium, which was once used to refer to caustic soda. In Latin, the name of the atomic no of sodium is found to be atrium. Potassium's symbol is K, which is derived from the Latin word kalium, which denoted alkali or potash. An element symbol's initial letter is capitalised. It's lowercase when there's a second letter.

Examination of 1 to 20 elements of the periodic table

Hydrogen

NEET Highest Scoring Chapters & Topics

This ebook serves as a valuable study guide for NEET exams, specifically designed to assist students in light of recent changes and the removal of certain topics from the NEET exam.

Download EBook

Under normal circumstances, hydrogen is a nonmetallic, colourless gas. It transforms into an alkali metal under great pressure. This element has three isotopes, each with a different amount of neutrons in its atoms. Protium is the most prevalent isotope. Deuterium and tritium are the other two.

The atomic number of Hydrogen is 1
H is its symbol
It is estimated to have an atomic mass value of 1.008 amu
H is assigned 1s¹ configuration
It is a Nonmetal belonging to group 1 and s-block

2. Helium

Helium is a light gas with no visible colour.
The atomic number of Helium is 2
He is its symbol
It is estimated to have an atomic mass value of 4.002 amu
He is assigned 1s² configuration
It is belonging to group 18 and s-block

3. Lithium

It is a highly reactive solid metal with a silver colour

The atomic number of Lithium is 3
Li is its symbol
It is estimated to have an atomic mass value of 6.94 amu
Li is assigned [He] 2s¹ configuration
It is an alkali metal belonging to group 1 and s-block.

4. Beryllium

It is a solid material with a shiny grey- white appearance

The atomic number of beryllium is 4
Be is its symbol
It is estimated to have an atomic mass value of 9.012 amu
Be is assigned [He] 2s² configuration
It is an alkaline earth metal belonging to group 2 and s-block.

5. Boron

The atomic number of Boron is 5
B is its symbol
It is estimated to have an atomic mass value of 10.81 amu
B is assigned [He] 2s² 2p¹ configuration
It is a metalloid belonging to group 13 and p-block.

6. Carbon

The atomic number of carbon is 6
C is its symbol
It is estimated to have an atomic mass value of 12.011 amu
C is assigned [He] 2s² 2p² configuration
It is a metalloid belonging to group 14 and p-block

7. Nitrogen

the atomic number of nitrogen is 7
N is its symbol
It is estimated to have an atomic mass value of 14.007 amu
N is assigned [He] 2s² 2p³ configuration
It belongs to group 15 and p-block.

8. Oxygen

The atomic number of oxygen is 8
O is its symbol
It is estimated to have an atomic mass value of 16 amu
O is assigned [He] 2s² 2p⁴configuration
It is belonging to group 16 and p-block

9. Fluorine

The atomic number of fluorine is 9
F is its symbol
It is estimated to have an atomic mass value of 18.989 amu
F is assigned [He] 2s² 2p⁵ configuration
It is belonging to group 17 and p-block

10. Neon

the atomic number of neon is 10
Ne is its symbol
It is estimated to have an atomic mass value of 20.179 amu
N is assigned [He] 2s² 2p⁶ configuration
It is belonging to group 18 and p-block

11. Sodium

The atomic number of sodium is 11
Na is the symbol of sodium
The Latin name of sodium is Natrium
It is estimated to have an atomic mass value of 22.989 amu
F is assigned [Ne] 3s¹ configuration
It is belonging to group 1 and s-block

12. Magnesium

The atomic number of magnesium is 12
Mg is its symbol
It is estimated to have an atomic mass value of 24.309 amu
Mg is assigned [Ne] 3s² configuration
It is belonging to group 2 and s-block

13. Aluminium

The atomic number of aluminium is 13
Al is its symbol
It is estimated to have an atomic mass value of 26.968 amu
Al is assigned [Ne] 3s² 3p¹ configuration
It is belonging to group 13 and p-block

14. Silicon

The atomic number of Silicon is 14
Si is its symbol
It is estimated to have an atomic mass value of 28.085 amu
Si is assigned [Ne] 3s² 3p² configuration
It belongs to group 14 and p-block.

15. Phosphorous

The atomic number of Phosphorous is 15
P is its symbol
It is estimated to have an atomic mass value of 30.9737 amu
P is assigned [Ne] 3s² 3p³ configuration
It is belonging to group 15 and p-block

16. Sulphur

The atomic number of Sulfer is 16
S is its symbol
It is estimated to have an atomic mass value of 32.09 amu
S is assigned [Ne] 3s² 3p⁴ configuration
It is belonging to group 16 and p-block

17. Chlorine

The atomic number of Chlorine is 17.
Cl is its symbol
It is estimated to have an atomic mass value of 35.45 amu
Cl is assigned [Ne] 3s²3p⁵ configuration
It is belonging to group 17 and p-block

18. Argon

the atomic number of argon is 18
Ar is its symbol
It is estimated to have an atomic mass value of 39.948 amu
Ar is assigned [Ne] 3s² 3p⁶ configuration
It is belonging to group 18 and p-block

19. Potassium

The atomic number of Potassium is 19
The Latin name of potassium is found to be Kalium
K is its symbol
It is estimated to have an atomic mass value of 39.093 amu
K is assigned [Ar] 4s¹ configuration
It is belonging to group 1 and s-block

20. Calcium

The atomic number of Calcium is 20.
Ca is its symbol
It is estimated to have an atomic mass value of 40.10 amu
Ca is assigned [Ar] 4s² configuration
It is belonging to group 14 and s-block

Also Read:

Development Of Modern Periodic Table	Mendeleev’s Periodic Table
Physical And Chemical Properties Of Elements	Nomenclature Of Elements With Atomic Number
Electronic Configuration In Periods And Groups	Atomic Size & Atomic Radius
Classification Of Elements And Periodicity In Properties	Ionisation Enthalpy

Some Solved Examples

Example.1 Choose the correct option:

1) (correct)The period of the element is determined by its highest shell

2)The period of the element is determined by its last orbital

3)The period of the element is determined by its valence shell electrons

4)The period of the element is determined by its valency

Solution

The period of the element is determined by its highest shell.

Hence, the answer is the option (1).

Example.2 Which pair of atomic numbers represents s-block elements

1)7,15

2)6,12

3)9,17

4) (correct)3,12

Solution

Out of the given elements, Z= 3 (Li) and Z= 12 (Mg) belong to the s Block of the periodic table.

Hence, the answer is the option (4).

Example.3 Newland’s octave law was successful in arranging:

1)Heavier elements

2) (correct)Lighter elements

3)Both

4)None

Solution

Newland’s octave law was successful in arranging lighter elements. After calcium, this law did not work accordingly.

Hence, the answer is the option (2).

Summary

The first 20 elements of the periodic table possess a range of chemical properties and their uses from simplest to most abundant elements like hydrogen and some essential elements of earth such as carbon, nitrogen and hydrogen these three elements are the main building blocks of many processes and also plays a vital role in the chemistry field and has various technological applications. Their characteristics are very diverse and interconnected. The first 20 elements of the periodic table contain metals such as sodium, magnesium, and aluminium. Metalliods such as silicon, and boron. Also, contains noble gases like neon and argon. Understanding all these elements helps us to appreciate the balance of nature and their application in chemistry.

Frequently Asked Questions (FAQs)

1. What are the first 20 elements of the periodic table?

The first 20 elements of the periodic table are:
1. Hydrogen (H)
2. Helium (He)
3. Lithium (Li)
4. Beryllium (Be)
5. Boron (B)
6. Carbon (C)
7. Nitrogen (N)
8. Oxygen (O)
9. Fluorine (F)
10. Neon (Ne)
11. Sodium (Na)
12. Magnesium (Mg)
13. Aluminum (Al)

14. Silicon (Si)
15. Phosphorus (P)
16. Sulfur (S)
17. Chlorine (Cl)
18. Argon (Ar)
19. Potassium (K)
20. Calcium (Ca)
These elements represent a range of different properties and are fundamental to many aspects of chemistry.

2. Why do elements have symbols, and what do they represent?

Each chemical element has a symbol, which is usually derived from its English name or its Latin name. For example, hydrogen is represented by "H" and sodium by "Na," which comes from its Latin name "Natrium." Symbols are used to simplify chemical formulas, equations, and labels for easier communication in science. They help avoid confusion and allow scientists globally to share knowledge efficiently.

3. Are there specific groups or categories among the first 20 elements?

Yes, the first 20 elements can be categorized into several groups based on their properties. For example, the first two elements, hydrogen and helium, are gases at room temperature, while lithium, beryllium, sodium, and magnesium are metals. Additionally, carbon, nitrogen, oxygen, and sulfur are essential non-metals. The different categories reflect varying properties, such as reactivity and state of matter.

4. Why is it important to learn about the first 20 elements?

Learning about the first 20 elements is crucial because they form the basis of all matter. Understanding these elements helps students grasp fundamental concepts in chemistry, such as bonding, reactions, and the nature of compounds. Many of these elements are vital for life, such as carbon, nitrogen, and oxygen, which are essential components of biological molecules.

5. How do the properties of the first 20 elements influence their use in everyday life?

The properties of the first 20 elements greatly influence their applications in our daily lives. For instance, carbon is a key element in organic chemistry, forming the backbone of biological molecules like proteins and carbohydrates. Oxygen is essential for respiration, while metals like sodium and potassium play crucial roles in physiological functions in living organisms. Moreover, elements like silicon are used in technology, underpinning modern electronics. Understanding these properties can lead to better insights into their practical uses.

6. Why are the first 20 elements important in chemistry?

The first 20 elements are crucial in chemistry because they form the foundation of the periodic table and represent the most common elements in nature. They include all the major elements necessary for life (C, H, O, N) and many elements essential for various chemical reactions and industrial processes. Understanding these elements provides a solid basis for comprehending more complex chemical concepts and interactions.

7. How does the electron configuration change as we move from hydrogen to calcium?

As we move from hydrogen (atomic number 1) to calcium (atomic number 20), the electron configuration becomes more complex. Each element adds one proton and one electron, filling electron shells in a specific order: 1s, 2s, 2p, 3s, 3p, and 4s. This progression follows the Aufbau principle, where electrons occupy the lowest energy orbitals first. By calcium, we see the first instance of electrons in the fourth energy level (4s).

8. Why does fluorine have a higher electronegativity than oxygen?

Fluorine has a higher electronegativity than oxygen despite oxygen having more protons because electronegativity is influenced by both nuclear charge and atomic size. Fluorine's smaller atomic radius means its nucleus exerts a stronger pull on electrons in chemical bonds. Additionally, fluorine needs only one electron to complete its octet, while oxygen needs two, making fluorine more "eager" to attract electrons.

9. How do the first 20 elements demonstrate periodicity?

The first 20 elements clearly demonstrate periodicity - the repetition of chemical and physical properties at regular intervals. This is evident in trends such as atomic size, ionization energy, and electronegativity. For example, the alkali metals (Li, Na, K) show similar chemical properties, as do the halogens (F, Cl). This periodicity forms the basis for the organization of the periodic table.

10. Why does potassium have a lower first ionization energy than sodium?

Potassium has a lower first ionization energy than sodium despite being lower in the periodic table because of the electron shielding effect. Potassium's outer electron is in the 4s orbital, which is further from the nucleus and more shielded by inner electrons than sodium's 3s electron. This makes it easier to remove potassium's outer electron, resulting in a lower ionization energy.

11. Why does helium have a full outer shell with only two electrons?

Helium has a full outer shell with only two electrons because its first (and only) electron shell can hold a maximum of two electrons. This is unique to the first energy level (n=1). All other energy levels can hold more electrons. Helium's stable electron configuration contributes to its chemical inertness and its classification as a noble gas.

12. What's the significance of the octet rule for the first 20 elements?

The octet rule is significant for the first 20 elements because it explains their bonding behavior. This rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons (like the noble gases). However, it's important to note that while this rule applies to many elements, there are exceptions, such as hydrogen (which is stable with 2 electrons) and elements beyond the second period.

13. How does the atomic radius trend change across the first 20 elements?

Across the first 20 elements, the atomic radius generally decreases from left to right across a period and increases from top to bottom down a group. This trend is due to increasing nuclear charge across a period (pulling electrons closer) and the addition of new electron shells down a group. However, there are some variations, particularly with the transition metals.

14. How does the reactivity of Group 1 elements (alkali metals) change as we move down the group?

The reactivity of Group 1 elements (alkali metals) increases as we move down the group from lithium to potassium. This is because the atomic radius increases, making it easier for the outermost electron to be removed. Additionally, there's more shielding from inner electrons, weakening the nucleus's hold on the valence electron. As a result, potassium is more reactive than sodium, which is more reactive than lithium.

15. Why does neon not form compounds despite not having a full outer shell like other noble gases?

Neon doesn't form compounds despite not having a full outer shell (like argon or krypton) because its electron configuration is still extremely stable. Neon has a complete 2p subshell, which provides a level of stability similar to a full outer shell. The energy required to add or remove electrons from neon is very high, making it energetically unfavorable to form compounds under normal conditions.

16. How does the concept of isoelectronic species apply to the first 20 elements?

Isoelectronic species are atoms or ions with the same number of electrons. Among the first 20 elements, we can find several examples of isoelectronic species. For instance, Ne, F⁻, O²⁻, N³⁻, and Na⁺ are all isoelectronic with 10 electrons each. Understanding isoelectronic species helps in predicting similar properties and behaviors among different atoms and ions, especially in terms of size and reactivity.

17. Why does nitrogen typically form three bonds instead of five, despite having five valence electrons?

Nitrogen typically forms three bonds instead of five because of its small atomic size and high electronegativity. While nitrogen has five valence electrons, it usually shares only three of them to complete its octet. Forming five bonds would require expanding its octet, which is energetically unfavorable for small, highly electronegative atoms like nitrogen. The triple bond in N₂ is exceptionally strong, contributing to nitrogen's stability in its molecular form.

18. How does the metallic character change across the first 20 elements?

Metallic character generally increases from right to left across a period and from top to bottom down a group among the first 20 elements. This is because metallic character is related to an element's tendency to lose electrons. Elements on the left side of the periodic table (like sodium and potassium) are more metallic, while those on the right (like chlorine and neon) are non-metallic.

19. How does the concept of electron affinity apply to the halogens in the first 20 elements?

Electron affinity is particularly relevant to the halogens (fluorine and chlorine) in the first 20 elements. These elements have high electron affinities, meaning they readily accept an electron to form negative ions. This is due to their electronic configuration being one electron short of a stable noble gas configuration. Fluorine has the highest electron affinity of all elements, making it extremely reactive. This property contributes to the halogens' strong oxidizing nature and their tendency to form ionic compounds with metals.

20. Why does the atomic mass not always increase consistently with atomic number in the first 20 elements?

The atomic mass doesn't always increase consistently with atomic number in the first 20 elements due to the presence of isotopes and the way average atomic mass is calculated. Some elements have multiple naturally occurring isotopes with different numbers of neutrons. The atomic mass we see on the periodic table is a weighted average of these isotopes based on their natural abundance. This can lead to situations where an element with a higher atomic number has a lower average atomic mass than the element preceding it.

21. Why does the melting point trend across period 3 (from sodium to chlorine) show an irregular pattern?

The melting point trend across period 3 shows an irregular pattern due to changes in bonding types and strengths. Sodium and magnesium have metallic bonds, with magnesium's higher melting point due to stronger bonding. Aluminum continues this trend. However, silicon has strong covalent network bonding, causing a sharp increase in melting point. Phosphorus, sulfur, and chlorine form molecular structures with weaker intermolecular forces, resulting in lower melting points. This variation in bonding types causes the irregular melting point trend.

22. Why does beryllium have a much higher melting point than lithium?

Beryllium has a much higher melting point than lithium due to its stronger metallic bonding. Beryllium has two valence electrons compared to lithium's one, allowing it to form stronger bonds within its crystal structure. Additionally, beryllium's smaller atomic size leads to a higher charge density, further strengthening the metallic bonds. These factors contribute to beryllium's significantly higher melting point.

23. How does the concept of effective nuclear charge explain trends in the first 20 elements?

Effective nuclear charge explains many trends in the first 20 elements. It represents the net positive charge experienced by an electron from the nucleus, accounting for electron shielding. As we move across a period, the effective nuclear charge increases, leading to trends like decreasing atomic radius and increasing ionization energy. Down a group, the increase in nuclear charge is largely offset by additional electron shielding, resulting in more gradual changes in properties.

24. Why does boron often form electron-deficient compounds?

Boron often forms electron-deficient compounds because it has three valence electrons but needs eight to complete its octet. Unlike its neighbors carbon and nitrogen, boron cannot easily gain or share enough electrons to achieve a full octet in many of its compounds. This leads to the formation of compounds with fewer than eight electrons around boron, making them electron-deficient and often giving boron unique chemical properties.

25. How does the first ionization energy trend differ between main group and transition elements among the first 20 elements?

The first ionization energy trend differs between main group and transition elements among the first 20 elements. For main group elements, ionization energy generally increases across a period and decreases down a group. However, for the transition elements (scandium and titanium in this range), the trend is less pronounced. This is because transition elements remove electrons from their 4s orbital first, not their partially filled 3d orbitals, leading to smaller changes in ionization energy.

26. Why does carbon have such a wide variety of allotropes compared to other elements in the first 20?

Carbon has a wide variety of allotropes compared to other elements in the first 20 due to its unique ability to form strong covalent bonds with itself in different spatial arrangements. Carbon has four valence electrons and can form single, double, or triple bonds. This versatility allows carbon to create various structures like diamond (tetrahedral), graphite (layered sheets), and fullerenes (spherical or cylindrical). No other element in this range has this combination of bonding flexibility and strength.

27. How does the concept of diagonal relationships apply to the first 20 elements?

Diagonal relationships in the periodic table refer to similarities in properties between certain elements diagonally adjacent to each other. Among the first 20 elements, this is most notably seen between lithium and magnesium, and between beryllium and aluminum. These pairs show similarities in properties due to a balance between size and charge effects. For example, lithium and magnesium have similar charge-to-size ratios, leading to comparable chemical behaviors in some reactions.

28. Why does silicon form a network covalent structure while carbon can form discrete molecules?

Silicon forms a network covalent structure, unlike carbon, due to its larger atomic size and lower electronegativity. Silicon's larger size makes it energetically favorable to form single bonds with four other silicon atoms, creating a tetrahedral network similar to diamond. Carbon, being smaller, can form strong multiple bonds, allowing it to create discrete molecules like CO₂ or CH₄. This difference explains why silicon exists as a solid at room temperature while carbon can form gases and liquids.

29. How does the concept of shielding effect explain the larger atomic size of potassium compared to sodium?

The shielding effect explains potassium's larger atomic size compared to sodium. Potassium's outer electron is in the 4s orbital, while sodium's is in the 3s orbital. The additional inner electron shell in potassium (the 3p electrons) shields the outer electron from the nuclear charge more effectively than in sodium. This increased shielding reduces the effective nuclear charge experienced by the outer electron, allowing it to exist further from the nucleus, resulting in a larger atomic radius for potassium.

30. How does the concept of electronegativity explain the trend in bond polarity across the first 20 elements?

Electronegativity, the ability of an atom to attract shared electrons in a bond, explains bond polarity trends across the first 20 elements. Electronegativity generally increases from left to right across a period and decreases down a group. This means that bonds between elements on the left (like sodium) and right (like chlorine) of the periodic table tend to be more polar. The greater the electronegativity difference, the more polar the bond. This concept helps predict the nature of chemical bonds and the properties of compounds formed by these elements.

31. Why does argon have a higher boiling point than neon, despite both being noble gases?

Argon has a higher boiling point than neon, despite both being noble gases, due to stronger intermolecular forces. Although both are nonpolar, argon has more electrons and thus a larger electron cloud. This larger electron cloud is more easily polarized, leading to stronger instantaneous dipole-induced dipole forces (London dispersion forces) between argon atoms. These stronger intermolecular forces require more energy to overcome, resulting in a higher boiling point for argon compared to neon.

32. How does the concept of effective nuclear charge explain the smaller atomic radius of aluminum compared to sodium?

The concept of effective nuclear charge explains why aluminum has a smaller atomic radius than sodium. Although aluminum has more electrons, its outer electrons are in the same shell (n=3) as sodium's. However, aluminum has a higher nuclear charge (13 protons vs. 11 for sodium) and less effective shielding from its inner electrons. This results in a greater effective nuclear charge in aluminum, pulling its outer electrons closer to the nucleus and creating a smaller atomic radius despite having more electrons overall.

33. Why does phosphorus have multiple allotropes while nitrogen typically exists as diatomic molecules?

Phosphorus has multiple allotropes while nitrogen typically exists as diatomic molecules due to differences in their atomic size and bonding capabilities. Nitrogen, being smaller, forms very strong triple bonds in N₂ molecules, making this form highly stable. Phosphorus, being larger, can form single, double, or triple bonds with itself. This versatility allows phosphorus to create various structural arrangements (allotropes) like white, red, and black phosphorus, each with different properties and stabilities.

34. How does the trend in electron affinity differ from the trend in electronegativity across the first 20 elements?

While electron affinity and electronegativity both generally increase from left to right across a period, they differ in their group trends. Electronegativity typically decreases down a group, but electron affinity doesn't follow a consistent trend down groups. For instance, chlorine has a higher electron affinity than fluorine, despite fluorine being more electronegative. This difference arises because electron affinity is more influenced by the specific electronic structure of each element, while electronegativity is more directly related to nuclear charge and atomic size.

35. Why does magnesium form a stronger ionic bond with oxygen than with fluorine, despite fluorine's higher electronegativity?

Magnesium forms a stronger ionic bond with oxygen than with fluorine, despite fluorine's higher electronegativity, due to the charge difference in the resulting ions. When magnesium bonds with oxygen, it forms Mg²⁺ and O²⁻ ions, resulting in a 2:2 charge ratio. With fluorine, it forms Mg²⁺ and F⁻ ions, a 2:1 charge ratio. The higher charges in the Mg-O bond lead to stronger electrostatic attraction. Additionally, the smaller size of the O²⁻ ion compared to two F⁻ ions allows for closer packing and stronger ionic interactions in magnesium oxide.

36. How does the concept of ionization energy explain the formation of different oxidation states in transition metals like scandium and titanium?

Ionization energy explains the formation of different oxidation states in transition metals like scandium and titanium. These elements can lose multiple electrons because the energy required to remove subsequent electrons (2nd, 3rd ionization energies) is not significantly higher than the first. This is due to the similar energies of their 4s and 3d orbitals. For example, titanium can form Ti²⁺, Ti³⁺, and Ti⁴⁺ ions because the energy required to remove these electrons is chemically accessible under various conditions, allowing for multiple oxidation states.

37. Why does beryllium have some properties similar to aluminum, despite being in a different group?

Beryllium shows some properties similar to aluminum due to their diagonal relationship in the periodic table. Both have a charge-to-size ratio that leads to similar chemical behaviors. Beryllium, like aluminum, forms amphoteric oxides and hydroxides, meaning they can react as both acids and bases. This similarity arises because beryllium's small size and high charge density make its chemistry more similar to aluminum than to its group members like magnesium. This diagonal relationship is an important concept in understanding periodic trends beyond simple group and period patterns.