Periodic Table

Periodic Table Trends:

1. Atomic Radius

2. Ionization Energy

3. Electron Affinity

4. Electronegativity

5. Density

6. Melting and Boiling Points

Introduction

The periodic table is a systematic arrangement of chemical elements based on their atomic number, electron configuration, and recurring chemical properties. Organized into rows (periods 1-7) and columns (groups or families 1-18), it serves as a framework for understanding the behavior of elements and their compounds.

The table can be generally divided into metals, nonmetals, and metalloids , each having distinct properties.

Metals, found on the left and center, are typically shiny, conductive, and malleable. They can be further divided into more detailed categories.
Nonmetals, located on the right, tend to be brittle, poor conductors, and diverse in appearance.
Metalloids exhibit properties of both metals and nonmetals and lie along a diagonal "stair-step" line.

Each element is represented by a unique symbol, with its atomic number (protons in the nucleus) displayed above. Groups share similar chemical properties, such as reactivity, due to their valence electrons. For example, Group 1 elements ( alkali metals ) are highly reactive, while Group 18 ( noble gases ) are stable and unreactive.

The periodic table not only explains chemical behavior but also predicts the discovery of new elements. Its modern design stems from Dmitri Mendeleev ’s 1869 version, which arranged elements by increasing atomic mass. Today, the periodic table is an essential tool in chemistry, physics, biology, and materials science.

Colors and Categories

The periodic table can be further divided into several categories, with the elements in each category expressing similar properties, both physical and chemical. In the periodic table above, there are 10 categories. Each category is represented with a unique color (differs among sources).

Alkali Metals (AM) contain the elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr).

Alkali metals are highly reactive elements in Group 1 of the periodic table. They have one valence electron, making them eager to form compounds. Soft, shiny, and low in density , they react vigorously with water to produce hydrogen gas and hydroxides. Their reactivity increases as you move down the group. $$2Na + 2H_2O \rightarrow 2NaOH + H_2\uparrow$$

Alkaline Earth Metals (AEM) contain the elements beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).

Alkaline earth metals are elements in Group 2 of the periodic table. They have two valence electrons, making them less reactive than alkali metals but still highly reactive. These metals are shiny, silvery-white, and good conductors. They react with water and acids, with reactivity increasing as you move down the group. $$Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2\uparrow$$

Transition Metals contain the elements 21-30, 39-48, 72-80, and 104-112. Examples include iron (Fe), copper (Cu), silver (Ag), tungsten (W), gold (Au), and mercury (Hg).

Transition metals, found in Groups 3–12 of the periodic table, are versatile elements with partially filled d orbitals. They exhibit high melting points, conductivity, and strength. Known for forming colorful compounds , they can have multiple oxidation states. Transition metals are essential in catalysis, industrial applications, and biological processes .

Post-transition Metals contain the elements aluminum (Al), gallium (Ga), indium (In), tin (Sn), thallium (Tl), lead (Pb), bismuth (Bi), nihonium (Nh), flerovium (Fl), moscovium (Mc), and livermorium (Lv).

Post-transition metals, located between transition metals and metalloids, are soft, malleable, and have lower melting points compared to transition metals. These elements, such as tin, lead, and aluminum, exhibit mixed metallic and nonmetallic properties. They are widely used in alloys, construction, and electronics due to their versatility and conductivity.

Lanthanides contain the elements 57-71, which include lanthanum (La), cerium (Ce), europium (Eu), thulium (Tm), and lutetium (Lu).

Lanthanides, often called rare earth elements, are a series of 15 metallic elements from lanthanum to lutetium. Known for their high reactivity and silvery-white appearance, they exhibit unique magnetic, optical, and catalytic properties. Lanthanides are essential in modern technologies like lasers , magnets, and batteries, despite being chemically similar and hard to separate.

Actinides contain the elements 89-103, which include actinium (Ac), thorium (Th), uranium (U), plutonium (Pu), berkelium (Bk), and lawrencium (Lr).

Actinides are a series of 15 metallic elements from actinium to lawrencium. They are radioactive and highly reactive, with some occurring naturally (like uranium and thorium) and others being synthetic. Actinides are crucial in nuclear energy and weapons due to their ability to release energy through fission but require careful handling due to their radioactivity. $$^{235}U + \text{Neutron} \rightarrow\ ^{236}U \rightarrow\ ^{141}Ba +\ ^{92}Kr + 3 \text{ Neutrons}$$

Metalloids contain the elements boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), and polonium (Po).

Metalloids are elements with properties that fall between metals and nonmetals, making them semiconductors and useful in electronics. They can conduct electricity under specific conditions and often exhibit a mix of metallic luster and brittle, nonmetallic behavior. Common examples include boron, silicon, and arsenic, essential for materials science and modern technology.

Nonmetals contain the elements hydrogen (H), carbon (C), nitrogen (N), oxygen (O), phosphorous (P), sulfur (S), and selenium (Se).

Nonmetals are elements characterized by poor electrical and thermal conductivity, low luster, and a tendency to gain or share electrons during chemical reactions. Found in various states—solids, liquids, and gases—they include essential elements like carbon, oxygen, and nitrogen. Nonmetals are crucial for life, energy processes, and forming covalent compounds.

Halogens contain the elements fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Te).

Halogens are highly reactive nonmetals in Group 17 of the periodic table. Known for their diatomic nature, they readily form salts with metals and are powerful oxidizing agents. Halogens include elements like fluorine and chlorine, vital in industrial, medical, and everyday applications, such as disinfectants , medications, and chemical synthesis.

Noble Gases contain the elements helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), radon (Rn), and oganesson (Og).

Noble gases are inert elements in Group 18 of the periodic table. They are colorless, odorless, and highly stable due to their full valence electron shells. Found naturally as monatomic gases, they are used in lighting, welding, and cooling systems. Examples include helium, neon, and argon, with minimal chemical reactivity.

In addition to categorizing the elements with respect to physical and chemical properties, we can categorize the elements with respect to the group the elements are in. Elements in the same group hold the same number of valence electrons. Some group names are already covered above.

Alkali metals (AM) are elements that are in Group 1.
Alkaline earth metals (AEM) are elements that are in Group 2.
Pnictogens (PN) are elements that are in Group 15, which contains the elements nitrogen, phosphorous, arsenic, antimony, bismuth, and moscovium.
Chalcogens (CH) are elements that are in Group 16, which contains the elements oxygen, sulfur, selenium, tellurium, polonium, and livermorium.
Halogens (HA) are elements that are in Group 17.
Noble gases (NG) are elements that are in Group 18.

Element ID Cards

Each element has a variety of physical and chemical properties. In the periodic table above, the properties are divided spatially into the left, upper right, and lower right side.

Left side: Summary
Upper right side: Physical properties
Lower right side: Chemical properties

In the summary section, the element ID contains the following properties:

1. Atomic number

The atomic number is a fundamental property of an element and is defined as the number of protons in the nucleus of an atom. It uniquely identifies each element on the periodic table and determines the element's chemical properties and position. For instance, hydrogen has an atomic number of 1, meaning its nucleus contains one proton, while oxygen has an atomic number of 8, indicating eight protons.

The atomic number also dictates the number of electrons in a neutral atom, influencing how atoms bond and interact with others. Arranging elements by increasing atomic number led to the structure of the modern periodic table, where elements are organized into periods and groups based on their electron configurations and similar chemical behavior. This concept provides the foundation for understanding atomic structure, reactivity, and the relationships between elements.

2. Element name

The element name is the unique identifier for each chemical element on the periodic table, reflecting its distinct identity and characteristics. These names are often derived from various sources, including Latin or Greek words, the element's properties, mythology, geographical locations, or the names of scientists who made significant contributions. For example, "oxygen" comes from Greek roots meaning "acid-former," while "einsteinium" honors Albert Einstein.

Each element's name provides a universal language for scientists across the globe, facilitating clear communication and consistency in chemical research and applications. Elements are also assigned a one- or two-letter symbol, such as "H" for hydrogen or "Fe" for iron, which is used for shorthand notation in equations, diagrams, and chemical formulas.

3. Period and group number

Period and group numbers are key organizational features of the periodic table that provide insight into an element's properties and behavior.

Periods are the horizontal rows on the periodic table, numbered from 1 to 7. They represent the number of electron shells or energy levels an atom of an element has. As you move across a period, elements gain more protons and electrons, causing gradual changes in properties such as atomic size and reactivity. For example, elements on the left are metals, while those on the right are nonmetals.
Groups, also known as families, are the vertical columns, numbered from 1 to 18. Elements in the same group share similar chemical properties because they have the same number of valence electrons (electrons in the outermost shell). For instance, Group 1 elements are highly reactive metals, while Group 18 contains stable noble gases.

These numbers help predict reactivity, bonding behavior, and element trends, offering a framework for understanding chemical behavior systematically.

4. Atomic weight

Atomic weight, also known as relative atomic mass, is the weighted average mass of an element's naturally occurring isotopes, measured in atomic mass units (amu). The value is dimensionless and standardized based on the carbon-12 isotope, which is defined as exactly 12 amu. It reflects both the mass and relative abundance of each isotope in nature. Unlike atomic number, which is always a whole number, atomic weight is often a decimal because it accounts for isotope distribution.

For example, carbon's atomic weight is approximately 12.01 amu because its two most common isotopes, carbon-12 and carbon-13, exist in a specific natural ratio. We use atomic weight to determine molar masses in chemical reactions and understand elemental properties.

5. Block

The block refers to a classification of elements in the periodic table based on the type of atomic orbital in which their outermost electrons reside. The periodic table is divided into four primary blocks: s-block, p-block, d-block, and f-block. These blocks group elements with similar electron configurations and properties.

s-block includes Groups 1 and 2, along with hydrogen and helium. Elements here have their outermost electrons in an s orbital. They are highly reactive metals and include alkali and alkaline earth metals.
p-block consists of Groups 13 to 18, containing diverse elements like metals, nonmetals, and metalloids. Their outermost electrons occupy a p orbital.
d-block, also called the transition metals, spans Groups 3 to 12. These elements have their outermost electrons filling d orbitals, contributing to their unique properties like variable oxidation states and colorful compounds.
f-block includes the lanthanides and actinides. These elements have electrons filling f orbitals and are known for their complex electron configurations.

The block structure helps explain periodic trends, reactivity, and the unique behavior of elements.

6. Category

Category refers to the classification of elements based on their shared chemical and physical properties. These categories help organize the periodic table and explain trends in element behavior. Common categories are discussed above in "Colors and Categories".

7. Electron configuration

Electron Configuration describes the arrangement of electrons in an atom's electron shells and subshells. It explains how electrons occupy energy levels and orbitals around the nucleus, providing insights into an element’s chemical behavior and periodic properties.

The configuration follows the Aufbau Principle, which states that electrons fill the lowest energy orbitals first. It also adheres to the Pauli Exclusion Principle (no two electrons in an atom can have identical quantum numbers) and Hund’s Rule (electrons fill degenerate orbitals singly before pairing).

The notation uses numbers and letters, such as 1s² 2s² 2p⁶, where:

Numbers (1, 2, 3, etc.) represent principal energy levels (n).
Letters (s, p, d, f) indicate the type of orbital (shape of the electron cloud).
Superscripts (e.g., s² or d⁵) show the number of electrons in each subshell.

Electron configurations explain periodic trends, such as ionization energy, atomic radius, and reactivity. For example:

Noble gases have full outer shells, making them inert.
Alkali metals have one outermost electron, making them highly reactive.

This concept is essential for predicting how elements bond and interact in chemical reactions.

In the physical properties section, the element ID contains the following properties:

8. Description

The description of an element usually contains a brief summary of the element's physical and chemical properties, alongside some history of the element's discovery. For some elements, additional information about common compounds, uses, or isotopes might be included.

9. Density (at 295K)

Density $\rho$ is a physical property that measures how much mass is contained within a given volume of a substance. It is commonly expressed in units like grams per cubic centimeter (g/cm³) or kilograms per liter (g/L). The formula for density is: $$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$$ This means that the density of an object can be calculated by dividing its mass by its volume. Substances with higher density have more mass packed into a smaller volume, while those with lower density have less mass in the same space. Density can vary with temperature and pressure, especially for gases, since they expand or contract in response to these factors. In the periodic table, elements with similar densities often exhibit similar physical characteristics. For example, metals like lead and gold are known for their high density, whereas gases like hydrogen and oxygen have low density. Density plays a key role in determining whether an object will float or sink in a fluid.

10. Molar heat capacity (at 298K, 1bar)

Molar heat capacity is the amount of heat required to raise the temperature of one mole of a substance by one degree Celsius (or Kelvin). Molar heat capacity is typically expressed in units of joules per mole per degree per Kelvin (J/mol·K).

This property is essential for understanding how substances absorb and store heat. For example, substances with a high molar heat capacity can absorb a significant amount of heat without a large temperature change, while those with low molar heat capacity heat up quickly with less energy. Water, for instance, has a relatively high molar heat capacity, which is why it is effective in regulating temperature.

11. Melting and boiling points (at 1atm)

Melting and boiling points are physical properties of elements and compounds that signify phase changes. The melting point is the temperature at which a substance transitions from a solid to a liquid, while the boiling point is the temperature at which it transitions from a liquid to a gas. These values are influenced by the strength of intermolecular or atomic bonds within the substance.

For example, elements like tungsten have extremely high melting and boiling points due to strong metallic bonds, while noble gases like helium have very low values because of weak dispersion forces.

These properties help us identify substances and understand their behavior under different temperature conditions. They also play a significant role in industrial applications, material selection, and thermodynamic calculations. Elements with similar melting and boiling points often belong to the same group or category in the periodic table, reflecting shared bonding characteristics.

12. Heat of fusion and vaporization

The heat of fusion and heat of vaporization are thermodynamic properties that describe the energy required for phase changes. The heat of fusion is the amount of energy needed to convert a substance from solid to liquid at its melting point without changing its temperature. This energy breaks the bonds holding the solid structure together. For example, water’s heat of fusion is 334 J/g, which is the energy needed to melt ice into liquid water.

The heat of vaporization, on the other hand, refers to the energy required to change a liquid into a gas at its boiling point. It is typically much higher than the heat of fusion because breaking the intermolecular forces in a liquid to form a gas requires more energy. For water, the heat of vaporization is 2260 J/g.

Both properties help us in understanding phase transitions, material behavior, and energy transfer in processes like evaporation, condensation, and melting.

In the chemical properties section, the element ID contains the following properties:

13. First ionization energy

First ionization energy is the amount of energy required to remove the outermost (or most loosely bound) electron from a neutral atom in its gaseous state. This process creates a positively charged ion. Ionization energy is measured in units like kilojoules per mole (kJ/mol) or electron volts (eV).

Elements with low first ionization energy, such as alkali metals, lose electrons easily, making them highly reactive. Conversely, noble gases have very high ionization energies due to their stable electron configurations, making them chemically inert.

Ionization energy increases across a period (left to right) in the periodic table because the increasing nuclear charge pulls electrons closer to the nucleus, requiring more energy to remove them. It decreases down a group (top to bottom) as electrons are farther from the nucleus, reducing the attractive force.

14. Electron affinity

Electron affinity is the amount of energy released (or absorbed) when a neutral atom in the gaseous state gains an electron to form a negatively charged ion. It reflects an atom’s tendency to attract additional electrons and is measured in kilojoules per mole (kJ/mol) or electron volts (eV).

Atoms with high electron affinity, such as halogens, release significant energy when they gain an electron, making them highly reactive and eager to form negative ions. Conversely, elements like noble gases have very low or even positive electron affinity because their stable electron configurations make it energetically unfavorable to gain an electron.

Electron affinity generally increases across a period (left to right) in the periodic table as the nuclear charge increases, drawing electrons closer. It decreases down a group (top to bottom) because the added electron is further from the nucleus, reducing the attraction.

15. Electronegativity

Electronegativity is a measure of an atom’s ability to attract shared electrons in a chemical bond. It plays a crucial role in determining bond polarity and molecular interactions.

Electronegativity generally increases across a period (left to right) in the periodic table due to increasing nuclear charge, which pulls electrons closer. It decreases down a group (top to bottom) as atomic size increases, making it harder for the nucleus to attract bonding electrons.

Elements with high electronegativity, like oxygen and nitrogen, tend to form polar covalent bonds, while those with low electronegativity, like alkali metals, form ionic bonds with nonmetals. Electronegativity differences between atoms help determine bond types—nonpolar covalent, polar covalent, or ionic—which influence molecular properties such as solubility, melting points, and reactivity.

16. Atomic radius

Atomic radius refers to the size of an atom, typically measured as the distance from the nucleus to the outermost electron. It is an important property that influences an element’s chemical and physical behavior. Atomic radius is affected by the number of electron shells and the effective nuclear charge (the pull of protons on electrons).

In the periodic table, atomic radius decreases across a period (left to right) because the increasing nuclear charge pulls electrons closer to the nucleus. Conversely, atomic radius increases down a group (top to bottom) as additional electron shells are added, making the atom larger.

For example, hydrogen has a smaller atomic radius than lithium because it has fewer electron shells. Similarly, chlorine has a smaller atomic radius than sodium due to stronger nuclear attraction.

Atomic radius is key in explaining trends like ionization energy, electronegativity, and bond lengths in molecules.

17. Isotopes

Isotopes are variants of the same chemical element that have the same number of protons (and therefore the same atomic number) but differ in the number of neutrons within their nuclei. This difference in neutron count gives isotopes different mass numbers, but they retain identical chemical properties because their electron configurations remain the same.

For example, carbon has two common isotopes: carbon-12 and carbon-14. Carbon-12 has 6 protons and 6 neutrons, while carbon-14 has 6 protons and 8 neutrons. While both isotopes behave identically in chemical reactions, their differing masses influence their physical properties and stability.

Some isotopes are stable, while others are radioactive, meaning they decay over time and emit radiation. Radioactive isotopes, such as uranium-235 or iodine-131, are widely used in applications like nuclear energy, medical imaging, and cancer treatments. Isotopes also play a role in scientific fields such as radiometric dating and tracing biological processes.

18. Oxidation states

Oxidation states, also known as oxidation numbers, represent the degree of oxidation or loss of electrons for an atom in a chemical compound. They are expressed as mostly integers, which can be positive, negative, or zero, depending on whether the atom has lost, gained, or shared electrons. Oxidation states help describe how electrons are distributed in a molecule and help us understand redox reactions.

For example, in water (H₂O), hydrogen has an oxidation state of +1, and oxygen has an oxidation state of -2. In ionic compounds like NaCl, sodium has an oxidation state of +1, and chlorine is -1.

Many elements can exhibit multiple oxidation states depending on the compound. Transition metals, for instance, are known for their variable oxidation states due to their partially filled d orbitals.

Periodic Table Trends

Below the periodic table, there are 6 graphs. These graphs demonstrate periodic table trends. Those periodic table trends contain information that help us predict atomic behavior, as well as molecular behavior.

1. Atomic radius

There are 3 factors influencing atomic radius:

Effective nuclear charge
Electron shell
Electron repulsion

Across a period (left to right), atomic radius decreases. This occurs because as protons are added to the nucleus, the effective nuclear charge increases, pulling electrons closer without significant shielding from inner electrons. Despite more electrons being present, they are added to the same energy level, leading to a stronger attraction and a smaller atomic size.

Down a group (top to bottom), atomic radius increases. This happens because new electron shells are added, increasing the distance between the nucleus and the outermost electrons. Although nuclear charge also increases, the shielding effect from inner electron shells reduces its pull, allowing the outer electrons to be farther away.

Atomic radius contributes to a lot of periodic table trends listed below.

2. First ionization energy

There are 2 factors influencing ionization energy:

Effective nuclear charge
Atomic radius

Across a period (left to right), ionization energy increases. As more protons are added to the nucleus, the effective nuclear charge increases, pulling electrons closer and making them harder to remove. Additionally, electrons are added to the same energy level, so shielding remains constant, further strengthening nuclear attraction.

Down a group (top to bottom), ionization energy decreases. This occurs because new electron shells are added, increasing the distance between the nucleus and the outermost electron. The shielding effect from inner electron shells weakens the nuclear pull, making it easier to remove an electron.

These trends explain why metals, which have low ionization energies, readily lose electrons to form positive ions, while nonmetals, with high ionization energies, tend to gain electrons in chemical reactions.

3. Electron affinity

There are 2 factors influencing electron affinity:

Effective nuclear charge
Atomic radius

Across a period (left to right), electron affinity generally increases. Atoms on the right side, particularly halogens, strongly attract additional electrons to complete their valence shells, releasing more energy in the process. However, noble gases are an exception, as they have full outer shells and do not readily gain electrons.

Down a group (top to bottom), electron affinity generally decreases. As atomic size increases, the added electron is farther from the nucleus and experiences weaker attraction, resulting in less energy release.

There are some exceptions due to electron repulsions and subshell stability, such as the noble gases, alkaline earth metals, and nitrogen group elements (pnictogens), which have relatively low or even positive electron affinities. These trends help explain why nonmetals, especially halogens, readily gain electrons, while metals generally do not.

4. Electronegativity

There are 2 factors influencing electronegativity:

Effective nuclear charge
Atomic radius

Across a period (left to right), electronegativity increases. As protons are added to the nucleus, the effective nuclear charge increases, pulling electrons closer and making atoms more likely to attract bonding electrons. Nonmetals, especially fluorine, oxygen, and nitrogen, have the highest electronegativity values because they strongly attract electrons to complete their valence shells.

Down a group (top to bottom), electronegativity decreases. As atomic size increases, the valence electrons are farther from the nucleus and experience greater shielding from inner electrons. This reduces the nucleus’s ability to attract bonding electrons, making larger atoms less electronegative.

The most electronegative element is fluorine (3.98 on the Pauling scale), while metals, especially alkali and alkaline earth metals, have the lowest values. Noble gases generally do not have electronegativity values because they rarely form bonds.

5. Density

Density trends on the periodic table are complex because density depends on:

Atomic mass
Atomic radius (thus, atomic volume)

Across a period (left to right), density generally increases from Groups 1-6, then decreases toward the noble gases. Metals on the left (alkali and alkaline earth metals) have low densities, while transition metals in the center are denser due to their compact atomic structures and strong metallic bonding. Nonmetals on the right tend to have lower densities, with noble gases being the least dense.

Down a group (top to bottom), density generally increases. As atomic mass increases, atoms become heavier, and even though atomic size also increases, the added mass tends to outweigh the volume increase, leading to a higher density. This trend is especially evident in alkali and alkaline earth metals.

There are exceptions, particularly among transition metals, where electron configurations influence packing efficiency and atomic spacing.

6. Melting and boiling points

Trends on melting and boiling points are influenced by several factors, including bonding, atomic radius, and electron configuration.

Across a period (left to right), melting and boiling points generally increase, peak near the center (transition metals), then decrease. Metals on the left (alkali and alkaline earth metals) have relatively low melting and boiling points due to weaker metallic bonding. Transition metals have higher values because of strong metallic bonds and dense electron clouds. Nonmetals, particularly noble gases, have very low melting and boiling points since they exist as gases at room temperature.

Down a group (top to bottom), trends vary by element type. In metals, melting and boiling points generally decrease as atoms become larger and metallic bonds weaken. In nonmetals, melting and boiling points tend to increase due to stronger van der Waals forces in larger molecules.

There are a lot of exceptions, such as carbon (diamond), which has an extremely high melting point due to its strong covalent bonds.

Learning periodic table trends helps us understand how elements behave chemically and physically. These trends allow us to predict reactivity, bond formation, and element properties. They also explain why certain elements form specific compounds and why some are more reactive than others.

Where do they live? What do they look like?

Which of the following is a noble gas?
1. Oxygen
2. Nitrogen
3. Krypton
4. Fluorine
Which of the following is an alkali metal?
1. Magnesium
2. Potassium
3. Aluminum
4. Neon
Which period does calcium belong to?
1. 2
2. 3
3. 4
4. 5
Which element is located in Group 17 and commonly found in table salt?
1. Chlorine
2. Sulfur
3. Argon
4. Magnesium
Which of the following elements is liquid at room temperature?
1. Mercury
2. Lead
3. Carbon
4. Sodium
Which of these elements is classified as a metalloid?
1. Carbon
2. Silicon
3. Iron
4. Gold
Which element has the atomic number 26?
1. Copper
2. Zinc
3. Iron
4. Nickel
Which of the following elements is a halogen?
1. Oxygen
2. Fluorine
3. Neon
4. Calcium
Which of these elements is essential for making steel?
1. Copper
2. Iron
3. Gold
4. Helium
Which group contains the most reactive nonmetals?
1. 2
2. 13
3. 14
4. 15
Which period does calcium belong to?
1. Group 1 (Alkali Metals)
2. Group 2 (Alkaline Earth Metals)
3. Group 17 (Halogens)
4. Group 18 (Noble Gases)
Which element is known for being lightweight and used in aircraft construction?
1. Titanium
2. Lead
3. Aluminum
4. Uranium
Which of these elements is a lanthanide?
1. Thorium
2. Cerium
3. Radon
4. Chromium
Which of the following elements is a chalcogen?
1. Nitrogen
2. Oxygen
3. Magnesium
4. Rubidium
Which element is in Period 3 and is a metalloid?
1. Boron
2. Silicon
3. Phosphorus
4. Argon
Which of these elements is highly radioactive?
1. Neon
2. Radium
3. Calcium
4. Aluminum
Which of these elements is a gas at room temperature?
1. Bromine
2. Iodine
3. Neon
4. Lead
Which of these elements has the highest atomic number?
1. Nickel
2. Uranium
3. Cesium
4. Bromine
Which element is essential for bone strength and found in milk?
1. Sodium
2. Calcium
3. Iron
4. Zinc
Which of these elements is used in batteries and is highly reactive with water?
1. Lithium
2. Copper
3. Tin
4. Platinum

Answer

C B C A A B C B B D B C B B B B C B B A

How do they differ from each other?

Which of the following elements has the largest atomic radius?
1. Li
2. Na
3. K
4. Rb
Which of the following has the smallest atomic radius?
1. Cl
2. F
3. O
4. N
Which of the following has the highest ionization energy?
1. Na
2. Mg
3. Al
4. Si
Which element has the lowest ionization energy?
1. Be
2. B
3. C
4. Na
Which of the following has the highest electron affinity?
1. Cl
2. O
3. F
4. S
Which of the following elements has the largest ionization energy?
1. Li
2. Be
3. B
4. C
Which of the following ions has the largest radius?
1. Na⁺
2. Mg²⁺
3. F^–
4. O^2–
Which of the following elements has the highest electronegativity?
1. N
2. O
3. F
4. Cl
Which of the following has the smallest ionization energy?
1. K
2. Ca
3. Sc
4. Ti
Which of the following has the largest atomic radius?
1. S
2. O
3. Se
4. Te
Which of the following elements has the smallest atomic radius?
1. Na
2. Mg
3. Al
4. Si
Which of the following elements has the highest electron affinity?
1. N
2. O
3. F
4. B
Which of the following elements has the highest electronegativity?
1. P
2. Cl
3. I
4. Br
Which of the following ions has the smallest radius?
1. K⁺
2. Ca²⁺
3. Cl^–
4. F^–
Which of the following has the largest electronegativity?
1. Na
2. Mg
3. Al
4. F
Which of the following elements has the smallest electron affinity?
1. Li
2. N
3. B
4. C
Which of the following elements has the largest atomic radius?
1. Al
2. Si
3. P
4. S
Which of the following elements has the smallest ionization energy?
1. K
2. Ca
3. Al
4. Si
Which of the following ions has the smallest radius?
1. O^2–
2. F^–
3. Na⁺
4. Mg²⁺
Which of the following has the largest atomic radius?
1. I
2. Br
3. Cl
4. F

Answer

D B D D A D D C A D D C B B D A A A D A

Why does the atomic radius generally increase as you move down a group on the periodic table?

Why does ionization energy tend to increase as you move across a period from left to right on the periodic table?

Why do elements in the same group of the periodic table tend to share similar chemical properties?

Why does fluorine have a higher electronegativity than iodine, despite both being in the same group?

Why does the atomic radius decrease as you move from left to right across a period on the periodic table?

Why is it that noble gases have the highest ionization energies compared to other elements in their periods?

Why do elements like sodium and potassium have low ionization energies compared to elements like neon and argon?

Why do metals generally have lower electronegativity values than nonmetals?

Why does the electron affinity increase as you move from left to right across a period?

Why does the size of an anion (like Cl^–) tend to be larger than that of its neutral atom (Cl)?

What do they enjoy doing?

What is the electron configuration of nitrogen (N)?
1. 1s² 2s² 2p³
2. 1s² 2s² 2p⁴
3. 1s² 2s¹
4. 1s² 2s² 2p⁵
Which of the following elements has an electron configuration of 1s² 2s² 2p⁵?
1. Oxygen (O)
2. Fluorine (F)
3. Neon (Ne)
4. Chlorine (Cl)
What is the correct electron configuration of magnesium (Mg)?
1. 1s² 2s² 2p⁶ 3s²
2. 1s² 2s² 2p⁶ 3s¹
3. 1s² 2s² 2p⁶
4. 1s² 2s² 2p⁵
Which element has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁵?
1. Sulfur (S)
2. Chlorine (Cl)
3. Argon (Ar)
4. Phosphorus (P)
What is the electron configuration of the Ca²⁺ ion?
1. 1s² 2s² 2p⁶ 3s²
2. 1s² 2s² 2p⁶ 3s² 3p⁶
3. 1s² 2s² 2p⁶ 3s² 3p⁴
4. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Which of the following has the electron configuration [Ne] 3s² 3p¹?
1. Magnesium (Mg)
2. Silicon (Si)
3. Aluminum (Al)
4. Chlorine (Cl)
Which of the following has a noble gas electron configuration?
1. Na⁺
2. Cl
3. K
4. O
Which of the following elements has a partially filled d-subshell?
1. Neon (Ne)
2. Iron (Fe)
3. Sodium (Na)
4. Aluminum (Al)
Which of the following is the electron configuration of a neutral potassium (K) atom?
1. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
2. 1s² 2s² 2p⁶ 3s² 3p⁶
3. 1s² 2s² 2p⁶ 3s² 3p⁵
4. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Which of the following ions has an electron configuration of [Ar] 3d⁶?
1. Mn²⁺
2. Fe²⁺
3. Co²⁺
4. Ni²⁺
Which of the following elements has an electron configuration ending in 4s² 3d⁷?
1. Nickel (Ni)
2. Cobalt (Co)
3. Iron (Fe)
4. Manganese (Mn)
Which of the following has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹?
1. K
2. Na
3. Ca
4. Mg
Which of the following ions has the same electron configuration as neon?
1. F⁻
2. Na
3. Cl⁻
4. K⁺
Which of the following elements has the ground-state electron configuration [Ar] 4s¹ 3d¹⁰?
1. Copper (Cu)
2. Zinc (Zn)
3. Nickel (Ni)
4. Chromium (Cr)
Which of the following ions has an electron configuration of [Ar] 3d¹⁰?
1. Zn²⁺
2. Cu²⁺
3. Fe³⁺
4. Co²⁺

Answer

A B A B B C A B A B B A A A A

Why does the atomic radius of elements generally decrease across a period, and how does this relate to their ability to form covalent bonds?

Explain why fluorine (F) is more likely to form an anion rather than a cation, using its electron configuration as the basis for your reasoning.

How does the electron configuration of sodium (Na) explain why it forms a +1 ion and participates in ionic bonding?

Why does nitrogen (N) form three covalent bonds in compounds like ammonia (NH₃), and how does its electron configuration play a role in this bonding behavior?

Explain why the elements in Group 18 (noble gases) are generally unreactive and do not easily form bonds, based on their electron configurations.

Why does calcium (Ca) commonly exhibit an oxidation state of +2, and how does its electron configuration explain this tendency?

How does the electron configuration of carbon (C) allow it to form four covalent bonds, and why is this crucial for the structure of organic compounds like methane (CH₄)?

Considering their electron configurations, why are alkali metals (like lithium and sodium) highly reactive, and how do they typically bond with other elements?

Explain why oxygen (O) tends to form two covalent bonds and why this property is essential for the formation of water (H₂O). How does its electron configuration influence this bonding behavior?

How do the electron configurations of the halogens (like chlorine, Cl) explain why they commonly form -1 anions and participate in ionic bonding with metals?

Why do transition metals, like iron (Fe), often exhibit multiple oxidation states? Relate this to their electron configurations and their tendency to form complex ions in coordination compounds.

How does the electron configuration of hydrogen (H) influence its ability to bond with other elements, particularly in compounds like H₂O and H₂S?

Why does potassium (K) have a lower ionization energy than magnesium (Mg), and how does this difference in electron configuration affect their reactivity?

Considering their electron configurations, why do alkaline earth metals (like magnesium and calcium) often form +2 ions and bond with halogens to form ionic compounds?

Why is the electron configuration of chlorine (Cl) responsible for its ability to gain one electron and form a stable anion, and how does this influence its reactivity with metals like sodium (Na)?

¹H	Stable
²H	Stable
³H	12.32 year

⁴He	Stable
³He	Stable
⁵He	602.0 ysec

⁷Li	Stable
⁶Li	Stable
³Li	Stable

⁹Be	Stable
⁵Be	Stable
⁶Be	5.0 zsec

¹¹B	Stable
¹⁰B	Stable
⁶B	Stable

Periodic Table of Elements

Periodic Table Trends:

Navigating the Periodic Table

Getting to Know the Elements:

¹²C	Stable
¹³C	Stable
¹⁴C	5700 year

¹⁴N	Stable
¹⁵N	Stable
¹⁰N	143.0 ysec

¹⁶O	Stable
¹⁸O	Stable
¹⁷O	Stable

¹⁹F	Stable
¹³F	Stable
¹⁴F	500.0 ysec