Polymath Learning Centre Periodic Table Interactive App
Learn about Elements, Bonding, Reactivity & History (Sec 1 Level)
Element Categories (Hover to Highlight)
Element Categories Explained
Elements are grouped into categories based on their shared physical and chemical properties, which are related to their electron structure.
Alkali Metals (Group 1, excluding Hydrogen)
These are very reactive, soft, silvery metals with low densities. They easily lose their single valence electron to form +1 ions (like Na⁺ or K⁺). They react strongly with water and nonmetals. Examples: Lithium (Li), Sodium (Na), Potassium (K).
Alkaline Earth Metals (Group 2)
These metals are reactive (though generally less than alkali metals), harder, and denser than alkali metals. They lose their two valence electrons to form +2 ions (like Mg²⁺ or Ca²⁺). Examples: Beryllium (Be), Magnesium (Mg), Calcium (Ca).
Transition Metals (Groups 3-12)
This is a large group of typical metals: generally hard, strong, dense, with high melting points. They are good conductors of heat and electricity. Many form coloured compounds and can form ions with different positive charges (e.g., Iron can be Fe²⁺ or Fe³⁺). Examples: Iron (Fe), Copper (Cu), Gold (Au), Titanium (Ti).
Lanthanides
These are metallic elements with atomic numbers 57-71 (Lanthanum to Lutetium). They are often placed below the main table. They have similar chemical properties to each other and are sometimes called "rare earth elements" (though some are quite abundant). Used in magnets, lasers, and lighting. Examples: Lanthanum (La), Cerium (Ce), Neodymium (Nd).
Actinides
These are metallic elements with atomic numbers 89-103 (Actinium to Lawrencium), also placed below the main table. All actinides are radioactive. Some occur naturally (like Thorium and Uranium), while others are man-made. Used in nuclear reactors and weapons. Examples: Actinium (Ac), Thorium (Th), Uranium (U), Plutonium (Pu).
Post-Transition Metals
These are metals located to the right of the transition metals. They tend to be softer and have lower melting/boiling points than transition metals. Examples: Aluminum (Al), Gallium (Ga), Tin (Sn), Lead (Pb).
Metalloids
These elements have properties intermediate between metals and nonmetals. They are found along the "staircase" line separating metals and nonmetals. Many are semiconductors, meaning they conduct electricity under certain conditions, making them vital for electronics. Examples: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As).
Polyatomic Nonmetals
These nonmetals don't typically form simple diatomic molecules like O₂ or N₂ in their standard states. They often form more complex structures or networks. Examples: Carbon (C - exists as graphite, diamond), Phosphorus (P - exists as P₄), Sulfur (S - exists as S₈).
Diatomic Nonmetals
These nonmetals naturally exist as molecules containing two atoms (diatomic) under standard conditions. This category includes the Halogens (Group 17) as well as Hydrogen, Nitrogen, and Oxygen. They tend to gain or share electrons in reactions. Examples: Hydrogen (H₂), Nitrogen (N₂), Oxygen (O₂), Fluorine (F₂), Chlorine (Cl₂).
Noble Gases (Group 18)
These are very unreactive (inert) gases because they have a full outer shell of valence electrons, making them very stable. They exist as individual atoms. Examples: Helium (He), Neon (Ne), Argon (Ar).
Unknown Properties
These are typically very heavy, man-made (synthetic) elements at the bottom of the periodic table. They are highly radioactive and decay very quickly (short half-lives), making it extremely difficult to study their chemical properties. Their properties are often predicted based on trends rather than direct observation. Examples: Meitnerium (Mt, 109) to Oganesson (Og, 118).
Chemical Bonds: Holding Atoms Together
Why Do Atoms Bond? It's All About Electrons!
Atoms join together to become more stable. Think of the Noble Gases (like Helium, Neon, Argon) - they are very stable and unreactive because they have a full outer shell of electrons (also called the valence shell).
Most other atoms want to achieve this stable state! They do this by gaining, losing, or sharing electrons with other atoms. These electrons involved in bonding are the valence electrons - the ones in the outermost shell.
Let's look at Sodium (Na) as an example. It has 11 electrons arranged in shells. We write this arrangement, called the electron configuration, as 2, 8, 1:
Sodium has 2 electrons in the first shell, 8 in the second, and 1 valence electron in the third (outermost) shell (configuration 2, 8, 1). To become stable like Neon (2, 8), it's easiest for Sodium to lose that 1 valence electron. This is why it readily forms a Na⁺ ion and participates in ionic bonding.
Atoms like Chlorine (Cl: electron configuration 2, 8, 7) have 7 valence electrons. It's easier for them to gain 1 electron to get a full outer shell (like Argon: 2, 8, 8). This explains why Chlorine forms Cl⁻ ions or shares electrons in covalent bonds.
Understanding valence electrons is key to understanding the different types of chemical bonds:
1. Ionic Bonds (Give and Take Electrons)
Ionic bonds usually form between a metal (which tends to lose electrons) and a nonmetal (which tends to gain electrons).
- Metal atoms lose their valence electrons to become positively charged ions (cations).
- Nonmetal atoms gain electrons to fill their valence shell, becoming negatively charged ions (anions).
- The positive and negative ions are strongly attracted to each other. This attraction is the ionic bond.
Example: Sodium (Na, loses 1e⁻) gives 1 electron to Chlorine (Cl, needs 1e⁻).
Sodium becomes Na⁺ (cation), Chlorine becomes Cl⁻ (anion). They attract to form NaCl (table salt).
2. Covalent Bonds (Sharing Electrons)
Covalent bonds usually form between two nonmetals.
- Atoms share pairs of valence electrons so that both atoms effectively achieve a full outer shell.
- Sharing electrons creates a stable unit called a molecule.
Example: Two Chlorine atoms (Cl, each needs 1e⁻) share one pair of electrons.
Each Cl atom shares one electron, forming a single covalent bond (Cl-Cl) in a Cl₂ molecule.
Another Example: Water (H₂O). Oxygen shares electrons with two Hydrogen atoms.
3. Metallic Bonds (Sea of Electrons)
Metallic bonds occur between atoms in a metal element.
- Metal atoms pack closely together in a regular pattern (a lattice).
- Their valence electrons are delocalized - they don't belong to any single atom but can move freely throughout the metal structure, like a "sea" of electrons.
- This "sea" holds the positive metal ions (atoms that have lost their valence electrons) together.
- This explains why metals are good conductors of electricity (electrons can flow) and are malleable (can be hammered into shape - the ions can slide past each other without breaking the bond).
Positive metal ions (M⁺ circles) held together by a surrounding 'sea' of mobile valence electrons (e⁻ small circles).
Reactivity: How Elements Behave
Reactivity describes how likely an element is to take part in a chemical reaction. This depends heavily on its valence electrons (electrons in the outermost shell).
Why React? The Goal is Stability!
Most atoms want to achieve a stable electron configuration, like the Noble Gases (Group 18), which have a full outer shell. They react to gain, lose, or share electrons to reach this stable state.
- Noble Gases (Group 18): Already have a full outer shell. They are very stable and unreactive. Think Helium (He), Neon (Ne), Argon (Ar).
- Other Elements: Need to react to get a full outer shell.
Using the Periodic Table to Predict Reactivity
The position of an element in the Periodic Table tells us a lot about its reactivity:
- Groups (Columns): Elements in the same group have the same number of valence electrons and similar chemical properties.
- Periods (Rows): Going across a period, the number of valence electrons increases.
Reactivity Trends for Metals:
- Alkali Metals (Group 1): Have only 1 valence electron. They are very reactive because it's easy to lose that one electron to get a stable configuration. Reactivity increases as you go down the group (Potassium (K) is more reactive than Sodium (Na), which is more reactive than Lithium (Li)). Why? The outer electron is further from the nucleus and easier to lose.
- Alkaline Earth Metals (Group 2): Have 2 valence electrons. They are reactive (but generally less than Group 1) as they need to lose 2 electrons. Reactivity also increases down the group.
Reactivity Trends for Nonmetals:
- Halogens (Group 17): Have 7 valence electrons. They are very reactive because they only need to gain 1 electron to get a full outer shell. Reactivity decreases as you go down the group (Fluorine (F) is the most reactive nonmetal, Chlorine (Cl) is less reactive, Bromine (Br) even less). Why? Atoms get larger down the group, so the nucleus has less pull on incoming electrons.
- Other Nonmetals (Groups 14-16): Reactivity varies. They tend to gain or share electrons. Oxygen (Group 16) is quite reactive. Carbon (Group 14) forms the basis of many compounds by sharing electrons.
In summary: Elements closest to the Noble Gases (like Group 1 and Group 17) are generally the most reactive because they are just one step away from achieving a stable electron configuration.
How We Picture the Atom: Atomic Models
Our idea of what an atom looks like has changed over time as scientists discovered more! The image below shows some important steps in this journey:
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Key Early Models Explained
- Dalton's Model (~1803): John Dalton proposed that elements were made of tiny, indivisible spheres called atoms. Atoms of the same element were identical, and different elements had different atoms. (See first model in image).
- Thomson's Plum Pudding Model (~1904): After discovering the electron, J.J. Thomson suggested the atom was a sphere of positive charge with negative electrons scattered inside, like plums in a pudding. (See second model in image).
- Rutherford's Nuclear Model (~1911): Based on his gold foil experiment, Ernest Rutherford proposed a new model. He suggested the atom has a tiny, dense, positively charged nucleus at the center (containing positive protons and later found, neutral neutrons). He proposed electrons orbited this nucleus, and that most of the atom was empty space. (See third model in image). {/* Removed visible HTML comment */}
- Bohr's Model (~1913): Niels Bohr refined Rutherford's model. He proposed that electrons don't just orbit randomly, but exist in specific energy levels or shells at fixed distances from the nucleus, much like planets in specific orbits (though not exactly the same!). Electrons can jump between these levels by absorbing or releasing energy. This model helped explain why atoms emit specific colours of light when heated. (See fourth model in image).
These models show how scientific understanding evolves with new experiments and evidence. The modern model (Quantum Mechanical Model) is more complex, describing electrons in terms of probability clouds (orbitals) rather than fixed paths, but these earlier models are important steps in understanding the atom's structure.
History of the Periodic Table & Understanding Numbers
A Brief History
The Periodic Table wasn't created overnight! Scientists observed patterns in elements for many years.
- Early Attempts: Scientists like Döbereiner noticed groups of three elements (triads) with similar properties.
- Dmitri Mendeleev (1869): The "father" of the modern periodic table! He arranged the known elements (around 60 at the time) by increasing atomic mass. Crucially, he left gaps for undiscovered elements and predicted their properties based on the patterns he saw. His predictions turned out to be remarkably accurate!
- Henry Moseley (1913): Used X-rays to determine the atomic number (number of protons) for each element. He realized arranging elements by atomic number resolved some inconsistencies in Mendeleev's table. This is the basis of the table we use today.
- Modern Table: Includes elements discovered or synthesized since Moseley, arranged by atomic number, reflecting the underlying electron structure which determines chemical properties.
Understanding the Numbers
Each element tile shows key numbers:
- Atomic Number (Top Number): This is the most important number! It tells you the number of protons in the nucleus of an atom of that element.
- Each element has a unique atomic number (e.g., every Carbon atom has 6 protons).
- In a neutral atom, the atomic number also equals the number of electrons.
- Atomic Mass (Bottom Number on Table): This number represents the average mass of an atom of that element, considering its isotopes (atoms of the same element with different numbers of neutrons).
- It's approximately equal to the number of protons PLUS the number of neutrons in the nucleus (since protons and neutrons have most of the mass). This sum is often called the mass number for a specific isotope.
- The atomic mass on the table is an average, so it's often a decimal number. For simple calculations (like finding neutrons), you can often round the atomic mass to the nearest whole number to get the mass number of the most common isotope. The tiles on the main table show this rounded value.
- Number of Neutrons ≈ Rounded Atomic Mass (Mass Number) - Atomic Number
Example (Carbon-12): Atomic Number = 6 (6 protons). Common Mass Number (Rounded Atomic Mass) = 12. Number of Neutrons = 12 - 6 = 6 neutrons.
Note: The info box that appears when you click an element shows the more precise average atomic mass (decimal value).
Quick Quiz!
{/* Question 1 */}1. Which element has the electronic configuration 2.8.1?