The Periodic Table of Elements

• The periodic table organizes all 118 known natural elements in a specific order
• Mendeleev developed it in 1869, with updates based on new discoveries
• Elements are arranged by increasing atomic number (number of protons) in horizontal rows
• Rows restart when the last electron occupies the next energy level
• The table includes 18 vertical groups and 7 horizontal periods
• The first period has two elements: hydrogen and helium
• The second and third periods contain eight elements each, from lithium to neon and sodium to argon, respectively
• Periods from the fourth to seventh are "long periods" with 18 elements each, except for the last two having 32 each
• Groups are numbered with Arabic numerals (1-18) and Roman numerals (I-VIII), with the latter skipping transition elements (groups 3-12)

Element Group Properties

• Elements in the same group share similar chemical and physical properties due to their electron structure
• Group 1 elements (except hydrogen) are alkali metals.
• Group 2 elements are alkaline earth metals.
• Group 3 elements are earth metals
• Group 4 includes carbon
• Group 5 contains nitrogen
• Group 6 contains oxygen
• Group 7 is the halogens
• Group 8 consists of noble gases (helium [He], neon [Ne], argon [Ar], krypton [Kr], xenon [Xe], radon [Rn]), which are almost inert due to a stable external electron configuration

Lanthanides and Actinides

• The Lanthanides (atomic numbers 58-71) and Actinides (atomic numbers 90-103) are placed below the main table for readability but are transition metals in group 3

Element Categories

• Metals are lustrous, and conduct electricity and heat well, and are malleable
• Non-metals have poor conductivity
• Semimetals have intermediate properties
• Blocks in the periodic table: s (groups 1, 2, and helium), d (groups 3-12), p (groups 13-18 excluding helium), f (lanthanides and actinides), referencing the filling of energy sublevels

Octet Rule

• Gilbert Newton Lewis introduced the octet rule in 1916
• It describes chemical bond formation between atoms, mainly for main group elements (Roman numeral numbering)
• Atoms tend to gain or lose electrons to achieve a stable electron configuration with eight electrons in their valence shell
• A complete valence shell leads to energetic stability and reduced tendency to form bonds
• Exceptions include atoms with few electrons, excess electrons, or eight electrons

Shortage of Electrons

• Atoms like hydrogen, beryllium, and boron have too few valence electrons to reach an octet
• Hydrogen has one electron, beryllium two, and boron three
• These atoms do not achieve eight electrons regardless of bonding
• Borane (BH3) is an example where boron, with three electrons, cannot achieve an octet with three hydrogen atoms providing one electron each

Excess Electrons

• Atoms exceeding the octet rule are mostly in the third period or beyond
• Sulfur trioxide (SO3) is an example where sulfur and oxygen have six valence electrons each
• This leads to 12 electrons in the valence shell
• Phosphorus can also expand its octet

Noble Gases Behavior

• Helium, unlike other noble gases, has only two electrons in its valence shell, following a stable "duet" configuration
• Elements near helium in atomic number can also achieve stability with two valence electrons
• Elements generally aim for a stable electron configuration, becoming non-reactive or less reactive like noble gases

First Groups Elements

• Elements in the first groups tend to lose electrons via ionization, adopting the electron configuration of the preceding noble gas
• Elements in groups VI and VII tend to gain electrons, releasing energy (electron affinity) and achieving the configuration of the following noble gas

Atomic Radius

• Atomic radius measures the average distance between an atom's nucleus and the electron cloud
• It estimates the size, but electron positioning doesn't have a definite boundary at the quantum level
• Generally applicable to isolated atoms or those in molecules and depends on the type of chemical bond
• Atomic radius in a molecule can differ from that of an isolated atom due to bonding forces and interactions.

Functional Aspects

• Atoms consist of a nucleus (protons and neutrons) and orbiting electrons
• Electrons occupy orbitals, representing probability regions instead of defined paths
• Atomic radius relates to the probability of finding electrons in a space volume
• Being statistical, it's an average reflecting electron distribution

Influences on Atomic Radius

• Atom size depends on electron distribution, influenced by atomic number (protons in nucleus) and principal quantum number, which dictates electron energy levels

Atomic Radius Location

• Atomic radius is within atomic chemistry and quantum physics
• Each element has a specific atomic radius depending on proton count (nuclear charge) and electrons

Periodic Table Trends

• Atomic radius decreases across (left to right) a period - Increasing proton count strengthens nuclear charge, pulling electrons inward

• Atomic radius increases down a group (top to bottom) - Adding energy levels (more distant orbitals) extends average electron distance

Atomic Radius Calculation

• Methods vary with bond types.
  • Measuring isolated atoms using electron diffraction or photoelectron spectroscopy is complex but possible by analyzing electron distribution.
  • For covalently bonded atoms: measure the distance between nuclei and divide by two (e.g., hydrogen molecule (H2) has a distance of 0.074 nm, so the covalent radius is 0.037 nm)

Ionization Energy

• It's the energy required to remove an electron from a gaseous atom, turning it into a positive ion
• Stronger attraction between the positive nucleus and the negative electron means higher ionization energy

Ionization Units

• Measured in kilojoules per mole (kJ/mol) or electronvolts (eV).
• For example, hydrogen has a high ionization energy of about 13.6 eV because it has only one electron close to the nucleus

Ionization Increase Trends

• Along the periodic table it increases from left to right
• Atoms (e.g., fluorine, neon) have stronger nuclei with more protons and attract electrons more strongly, hence more energy is needed to remove
• Down the periodic table it decreases
• Electrons are easier to remove as they are farther from the nucleus

Ionization Examples

• Sodium (Na) easily loses an electron because it is on the left
• Neon (Ne) holds its electrons more tightly so requires more energy, since it is on the right
• Lithium (Li) holds its electrons strongly, having high energy requirements, due to it being at the top
• Cesium (Cs) has electrons far away and hence low ionization, since it is at the bottom

Electron Affinity

• It's the energy released when an atom gains an electron; negative if an anion forms, positive if a cation forms
• Atoms need energy to accept electrons
• Closer electrons have high affinity; more distant electrons have low affinity
• It increases across periods because atomic numbers increase, atoms shrink, and electrons get closer to the nucleus

Electron Affinity Decrease Trend

• Decreases down groups because nuclei grow, outer electrons separate, and energy lessens
• Exceptions exist as noble gases have low affinity as they are electron stable

Electronegativity

• Electronegativity measures an atom's ability to attract bonding electrons in molecules
• More electronegative atoms exert greater force on shared electrons
• It increases with nuclear charge and decreases as shell count rises Fluorine (F) is highly electronegative while lithium (Li) is weakly electro negative

Additional Details

• It increases from left to right and reduces from top to bottom
• It's on a scale like Pauling, and helps understand bonding tendencies