Properties and Data summary of Elements 1 to 30

Atomic Number	Symbol	Name	Protons	Neutrons	Atomic Weight (u)	Valency	Common Ions	State
1	H	Hydrogen	1	0	1.008	1	H⁺	Gas
2	He	Helium	2	2	4.0026	0	—	Gas
3	Li	Lithium	3	4	6.94	1	Li⁺	Solid
4	Be	Beryllium	4	5	9.0122	2	Be²⁺	Solid
5	B	Boron	5	6	10.81	3	B³⁺	Solid
6	C	Carbon	6	6	12.011	4	—	Solid
7	N	Nitrogen	7	7	14.007	3	N³⁻	Gas
8	O	Oxygen	8	8	15.999	2	O²⁻	Gas
9	F	Fluorine	9	10	18.998	1	F⁻	Gas
10	Ne	Neon	10	10	20.180	0	—	Gas
11	Na	Sodium	11	12	22.990	1	Na⁺	Solid
12	Mg	Magnesium	12	12	24.305	2	Mg²⁺	Solid
13	Al	Aluminum	13	14	26.982	3	Al³⁺	Solid
14	Si	Silicon	14	14	28.085	4	—	Solid
15	P	Phosphorus	15	16	30.974	3	P³⁻	Solid
16	S	Sulfur	16	16	32.06	2	S²⁻	Solid
17	Cl	Chlorine	17	18	35.45	1	Cl⁻	Gas
18	Ar	Argon	18	22	39.948	0	—	Gas
19	K	Potassium	19	20	39.098	1	K⁺	Solid
20	Ca	Calcium	20	20	40.078	2	Ca²⁺	Solid
21	Sc	Scandium	21	24	44.956	3	Sc³⁺	Solid
22	Ti	Titanium	22	26	47.867	4	Ti⁴⁺	Solid
23	V	Vanadium	23	28	50.942	5	V⁵⁺	Solid
24	Cr	Chromium	24	28	51.996	2	Cr³⁺, Cr⁶⁺	Solid
25	Mn	Manganese	25	30	54.938	7	Mn²⁺, Mn⁴⁺, Mn⁷⁺	Solid
26	Fe	Iron	26	30	55.845	2	Fe²⁺, Fe³⁺	Solid
27	Co	Cobalt	27	32	58.933	2	Co²⁺, Co³⁺	Solid
28	Ni	Nickel	28	31	58.693	2	Ni²⁺	Solid
29	Cu	Copper	29	35	63.546	2	Cu⁺, Cu²⁺	Solid
30	Zn	Zinc	30	35	65.38	2	Zn²⁺	Solid

Note: The number of neutrons is calculated based on the most common isotope of each element. Valency can vary depending on the chemical compound formed. The standard state refers to the physical state of the element at room temperature (298 K) and standard atmospheric pressure.

Metals, metalloids, and nonmetals. Below is a table that organizes these elements accordingly:

Category	Subcategory	Atomic Number	Symbol	Name
Metals	Alkali Metals	3	Li	Lithium
		11	Na	Sodium
	Alkaline Earth Metals	4	Be	Beryllium
		12	Mg	Magnesium
	Transition Metals	21	Sc	Scandium
		22	Ti	Titanium
		23	V	Vanadium
		24	Cr	Chromium
		25	Mn	Manganese
		26	Fe	Iron
		27	Co	Cobalt
		28	Ni	Nickel
		29	Cu	Copper
		30	Zn	Zinc
	Post-transition Metals	13	Al	Aluminum
		31	Ga	Gallium
Metalloids		5	B	Boron
		14	Si	Silicon
		32	Ge	Germanium
		33	As	Arsenic
Nonmetals	Reactive Nonmetals	1	H	Hydrogen
		6	C	Carbon
		7	N	Nitrogen
		8	O	Oxygen
		9	F	Fluorine
		15	P	Phosphorus
		16	S	Sulfur
		17	Cl	Chlorine
	Noble Gases	2	He	Helium
		10	Ne	Neon
		18	Ar	Argon

Note: Elements with atomic numbers 31 to 33 are included to complete the metalloid category, even though they are beyond the first 30 elements. The classification of elements, especially metalloids, can vary based on different sources. This table is based on standard periodic table categorizations.

Properties of Elements

A. METAL

i. Alkali Metals

Alkali metals are a group of chemical elements located in Group 1 of the periodic table. This group comprises lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). These elements share several characteristic properties due to their similar electron configurations.

Key Characteristics:

Physical Properties:

• Appearance: Alkali metals are soft, silvery-white metals that can be easily cut with a knife.
• Density: They have low densities, with lithium, sodium, and potassium being less dense than water.
• Melting and Boiling Points: These metals possess relatively low melting and boiling points compared to other metals, and these points decrease as one moves down the group.

Chemical Properties:

• Reactivity: Alkali metals are highly reactive, especially with water, forming hydroxides and releasing hydrogen gas. Their reactivity increases down the group, with lithium being the least reactive and francium the most.
• Electron Configuration: Each alkali metal has a single electron in its outermost shell, which it readily loses to form a cation with a +1 charge.
• Compounds Formation: They commonly form ionic compounds, such as halides (e.g., sodium chloride) and oxides.

Applications:

• Lithium: Used in rechargeable batteries, such as those in smartphones and electric vehicles.
• Sodium: Essential in the form of sodium chloride (table salt) and used in street lighting (sodium vapor lamps).
• Potassium: Vital for plant growth and commonly found in fertilizers.

Due to their high reactivity, especially with water and air, alkali metals are typically stored under oil or in inert atmospheres to prevent unwanted reactions.

ii. Alkaline Earth Metals

Alkaline earth metals are a group of six elements in Group 2 of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These elements share similar properties due to their comparable electronic configurations, each having two electrons in their outermost s-orbital.

Key Characteristics:

• Appearance: They are shiny, silvery-white metals.
• Reactivity: These metals are somewhat reactive at standard temperature and pressure, readily losing their two outermost electrons to form cations with a +2 charge.
• Density and Hardness: They have relatively low densities and are softer compared to transition metals.
• Melting and Boiling Points: They possess higher melting and boiling points than alkali metals but lower than most transition metals.

Chemical Properties:

• Formation of Compounds: Alkaline earth metals react with halogens to form ionic halides and with oxygen to form oxides. Except for beryllium, they react with water to form strongly alkaline hydroxides.

• Ionization Energies: They have relatively low first and second ionization energies, facilitating the loss of two electrons to achieve a stable electron configuration.

Applications:

• Magnesium and Calcium: Essential for biological processes; magnesium is a key component of chlorophyll, and calcium is vital for bones and teeth.
• Beryllium: Used in aerospace materials due to its stiffness, light weight, and dimensional stability over a wide temperature range.
• Strontium and Barium: Utilized in fireworks and flares for their characteristic colors; strontium imparts a crimson color, while barium produces a green hue.
• Radium: Historically used in luminescent paints, though its radioactivity has limited its modern applications.

iii. Transition Metals

Transition metals are elements found in the d-block of the periodic table, specifically in groups 3 to 12. They include elements such as iron (Fe), copper (Cu), and gold (Au), among others. These metals are characterized by their ability to form compounds with variable oxidation states and to form colored complexes.

Key Characteristics:

• Variable Oxidation States: Transition metals can exhibit multiple oxidation states, allowing them to form a variety of compounds.
• Formation of Colored Compounds: Many transition metal compounds are colored due to electronic transitions between d-orbitals.
• Magnetic Properties: Some transition metals and their compounds exhibit magnetic properties due to unpaired d-electrons.
• Catalytic Activity: They often serve as catalysts in chemical reactions, facilitating processes without being consumed.

Applications:

• Industrial Catalysts: Transition metals are widely used as catalysts in industrial processes, such as the use of iron in the Haber process for ammonia synthesis.
• Alloys: They are essential components in the production of various alloys, enhancing properties like strength and corrosion resistance.
• Biological Systems: Certain transition metals, like iron in hemoglobin, play crucial roles in biological systems.

iv. Post-Transition Metals

Post-transition metals, also known as poor metals, are a subset of metallic elements situated between the transition metals and metalloids in the periodic table. This category typically includes elements such as aluminum (Al), gallium (Ga), indium (In), tin (Sn), thallium (Tl), lead (Pb), and bismuth (Bi). These metals exhibit distinct physical and chemical properties that differentiate them from both transition metals and metalloids.

Physical Properties:

• Softness and Malleability: Post-transition metals are generally softer and more malleable compared to transition metals, making them easier to deform under pressure.
• Lower Melting Points: They possess relatively low melting points. For instance, gallium melts just above room temperature, at approximately 29.76°C (85.57°F).
• Poor Mechanical Strength: These metals often have lower tensile strength and hardness, limiting their use in applications requiring structural durability.

Chemical Properties:

• Higher Electronegativity: Post-transition metals typically have higher electronegativity values than transition metals, indicating a greater tendency to attract electrons in chemical bonds.
• Tendency to Form Covalent Bonds: Due to their position near the metal-nonmetal borderline, these metals often form covalent or directional bonds, especially in their crystalline structures.

Applications:

• Aluminum (Al): Widely used in the aerospace and construction industries due to its lightweight and corrosion-resistant properties.
• Tin (Sn): Commonly used as a protective coating for other metals and in the production of alloys like bronze and solder.
• Lead (Pb): Historically used in pipes and paints, now primarily utilized in lead-acid batteries and radiation shielding.

B. METALLOIDS

Metalloids are chemical elements that exhibit properties intermediate between those of metals and nonmetals, making them difficult to classify strictly as either. This unique positioning in the periodic table imparts a blend of characteristics from both categories.

Common Properties of Metalloids:

• Appearance: Metalloids typically have a metallic luster, giving them a shiny appearance similar to metals.
• Brittleness: Unlike metals, which are usually malleable and ductile, metalloids are often brittle and can break or shatter like nonmetals.
• Electrical Conductivity: They are generally semiconductors, meaning they have electrical conductivity between that of metals (conductors) and nonmetals (insulators). This property is temperature-dependent and forms the basis for their use in electronic components.
• Chemical Behavior: Metalloids can exhibit both metallic and nonmetallic chemical behaviors. For instance, they can form alloys with metals and also react with nonmetals to form covalent compounds.
• Oxide Formation: The oxides of metalloids are typically amphoteric or weakly acidic, contrasting with the strongly acidic oxides of nonmetals and the basic oxides of metals.

Elements Commonly Classified as Metalloids:

The classification of metalloids can vary, but the following elements are commonly recognized as metalloids:

1. Boron (B)
2. Silicon (Si)
3. Germanium (Ge)
4. Arsenic (As)
5. Antimony (Sb)
6. Tellurium (Te)
7. Polonium (Po)

These elements are typically found in a diagonal region of the p-block of the periodic table, extending from boron at the top left to polonium at the bottom right.

Applications of Metalloids:

Due to their semiconducting properties, metalloids, particularly silicon and germanium, are extensively used in the electronics industry for manufacturing transistors, diodes, and other semiconductor devices. Boron is utilized in glassmaking and as a dopant in semiconductors, while arsenic and antimony are used in alloys and certain types of semiconductors.

Understanding the unique properties of metalloids is essential in fields like materials science and electronics, where their intermediate characteristics are harnessed for various technological applications.

C. NON-METALS

Nonmetals are chemical elements that predominantly lack the characteristics of metals. They exhibit a wide range of physical and chemical properties, distinguishing them from metallic elements.

Physical Properties:

• States of Matter: Nonmetals can exist in various states at room temperature:
• Gases: Hydrogen (H₂), nitrogen (N₂), oxygen (O₂), fluorine (F₂), chlorine (Cl₂), and the noble gases (e.g., helium, neon).
• Liquids: Bromine (Br₂) is the only nonmetal that is liquid at room temperature.
• Solids: Carbon (C), phosphorus (P), sulfur (S), selenium (Se), and iodine (I₂) are solid nonmetals.

• Appearance: Solid nonmetals typically have a dull, non-lustrous appearance. For example, sulfur is a yellow, brittle solid, and carbon can exist as graphite, which is a soft, black material.
• Conductivity: Nonmetals are generally poor conductors of heat and electricity. This property makes them suitable as insulators in various applications.
• Density: They usually have lower densities compared to metals. For instance, the density of oxygen gas is significantly less than that of iron.
• Melting and Boiling Points: Nonmetals often have lower melting and boiling points than metals. However, there are exceptions; for example, carbon in the form of diamond has a very high melting point.

Chemical Properties:

• Electronegativity: Nonmetals typically have higher electronegativity values, meaning they have a greater tendency to attract electrons in chemical bonds. For example, fluorine is the most electronegative element.
• Ionization Energy: They possess higher ionization energies, making it more difficult to remove electrons from them. This contributes to their tendency to gain or share electrons during chemical reactions.
• Oxidation States: Nonmetals can exhibit a variety of oxidation states, often negative, reflecting their ability to gain electrons. For instance, sulfur can have oxidation states ranging from -2 to +6.
• Acidic Oxides: The oxides of nonmetals are generally acidic. When dissolved in water, they form acids. For example, sulfur dioxide (SO₂) dissolves in water to form sulfurous acid (H₂SO₃).

Examples of Nonmetals:

• Hydrogen (H): The lightest element, existing as a diatomic gas (H₂) under standard conditions.
• Carbon (C): Found in various allotropes, including graphite and diamond, each with distinct properties.
• Nitrogen (N): A diatomic gas (N₂) that makes up about 78% of Earth's atmosphere.
• Oxygen (O): A diatomic gas (O₂) essential for respiration in most living organisms.
• Phosphorus (P): Exists in several allotropes, with white phosphorus being highly reactive and red phosphorus being more stable.
• Sulfur (S): Commonly found as a yellow crystalline solid, used in the production of sulfuric acid.
• Halogens: This group includes fluorine (F), chlorine (Cl), bromine (Br), and iodine (I), all of which are reactive nonmetals.
• Noble Gases: Elements like helium (He), neon (Ne), and argon (Ar) are inert under most conditions due to their complete valence electron shells.

Applications of Nonmetals:

• Biological Importance: Elements like carbon, hydrogen, nitrogen, oxygen, phosphorus, and sulfur are fundamental to biological molecules and processes.
• Industrial Uses: Nonmetals are utilized in various industries. For example, nitrogen is used in the production of ammonia for fertilizers, and sulfur is used in vulcanizing rubber.
• Technological Applications: Silicon, though a metalloid, is often discussed alongside nonmetals due to its semiconductor properties, making it crucial in electronics.

i. Halogens

Halogens are a group of elements located in Group 17 (formerly Group VIIA) of the periodic table, comprising fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Ts). The term "halogen" originates from Greek, meaning "salt-former," reflecting their propensity to form salts when reacting with metals.

Physical Properties:

• States of Matter: Halogens exhibit all three classical states of matter at room temperature:

  • Gases: Fluorine and chlorine are diatomic gases.
  • Liquid: Bromine exists as a diatomic liquid.
  • Solid: Iodine and astatine are solid under standard conditions.

• Appearance: These elements display a range of colors:

  • Fluorine is a pale yellow gas.
  • Chlorine is a greenish-yellow gas.
  • Bromine is a reddish-brown liquid.
  • Iodine is a dark gray or purple-black solid.

Chemical Properties:

• Reactivity: Halogens are highly reactive, with reactivity decreasing down the group:
• Fluorine is the most reactive, capable of forming compounds with almost all elements, including some noble gases.
• Chlorine, bromine, and iodine follow in decreasing reactivity.
• Electronegativity: They possess high electronegativity values, with fluorine being the most electronegative element, leading to strong tendencies to attract electrons during chemical reactions.
• Diatomic Molecules: In their elemental form, halogens exist as diatomic molecules (e.g., F₂, Cl₂), meaning they form molecules consisting of two atoms.

Reactions with Water:

• Fluorine: Reacts vigorously with water to produce oxygen and hydrogen fluoride:

2F₂(g) + 2H₂O(l) → O₂(g) + 4HF(aq)

• Chlorine: Dissolves in water to form a mixture of hydrochloric acid and hypochlorous acid, which can act as a disinfectant or bleach:

Cl₂(g) + H₂O(l) ⇌ HCl(aq) + HClO(aq)

• Bromine: Slightly soluble in water, reacting slowly to form hydrogen bromide and hypobromous acid:

Br₂(l) + H₂O(l) ⇌ HBr(aq) + HBrO(aq)

• Iodine: Minimally soluble in water and does not react significantly; however, it can form soluble triiodide ions (I₃⁻) in the presence of iodide salts.

Applications:

• Fluorine: Used in the production of uranium hexafluoride for nuclear fuel processing and in manufacturing Teflon (polytetrafluoroethylene).
• Chlorine: Widely used as a disinfectant in water treatment, in the production of polyvinyl chloride (PVC) plastics, and as a bleaching agent.
• Bromine: Employed in flame retardants, certain types of photography chemicals, and as a fumigant.
• Iodine: Essential in medicine as an antiseptic and in the production of iodized salt to prevent iodine deficiency.

ii. Inert gases or Noble gases

Noble gases, also known as inert gases, are elements in Group 18 of the periodic table. They include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), radon (Rn), and the synthetic element oganesson (Og). These gases are characterized by their lack of chemical reactivity and share several key properties:

1. Chemical Inertness: Noble gases have complete valence electron shells, making them highly stable and largely unreactive under standard conditions. This full valence shell configuration means they have little tendency to gain or lose electrons, resulting in minimal chemical reactivity.
2. Monatomic Nature: Under standard conditions, noble gases exist as single atoms rather than molecules. This monatomic form is due to their stable electronic configuration, which negates the need to form bonds with other atoms.

3. Low Boiling and Melting Points: They have very low boiling and melting points, which increase with atomic number. For example, helium boils at 4.15 K, while xenon boils at 165.2 K. These low temperatures are attributed to weak intermolecular forces between the atoms.

4. Density: The density of noble gases increases with atomic number. Helium, the lightest, has a density of 0.1786 g/dm³ at 0°C and 1 atm, whereas radon, the heaviest natural noble gas, has a density of 9.97 g/dm³ under the same conditions.

5. Solubility: Noble gases are sparingly soluble in water, with solubility generally increasing with atomic number. For instance, helium has a solubility of 8.61 cm³/kg in water at 20°C, while xenon has a solubility of 108.1 cm³/kg.

6. Thermal Conductivity: They exhibit low thermal conductivity, which decreases with increasing atomic number. Helium has a thermal conductivity of 141.84 mW m⁻¹ K⁻¹ at 0°C and 1 atm, whereas xenon's is 5.06 mW m⁻¹ K⁻¹ under the same conditions.

7. Electronegativity and Ionization Energy: Noble gases have high ionization energies and negligible electronegativity values, reflecting their reluctance to form chemical bonds. Their ionization energies decrease with increasing atomic number.

These unique properties make noble gases suitable for various applications, such as providing inert atmospheres in industrial processes, filling gas-discharge lamps, and serving as cryogenic refrigerants.

This information is compiled from standard periodic table data and general chemical properties of the elements