🫴Physical Science Unit 5 – Periodic Table and Chemical Bonding

Key Concepts

• Atoms consist of protons, neutrons, and electrons with protons and neutrons in the nucleus and electrons orbiting in shells
• Elements arranged on the periodic table according to increasing atomic number (number of protons)
• Valence electrons in the outermost shell determine an element's chemical properties and bonding behavior
• Chemical bonds form when atoms share, donate, or accept electrons to achieve a stable electron configuration
• Ionic bonds involve the transfer of electrons between a metal and a nonmetal resulting in oppositely charged ions
• Covalent bonds involve the sharing of electrons between two nonmetals resulting in a shared electron pair
• Metallic bonds occur between metal atoms where valence electrons are delocalized and shared among all atoms
• Electronegativity, atomic radius, and ionization energy display periodic trends across the periodic table

Atomic Structure Basics

• Atoms are the fundamental building blocks of matter consisting of a dense nucleus surrounded by electrons in shells
• The nucleus contains positively charged protons and neutral neutrons while negatively charged electrons orbit the nucleus
• Protons and neutrons have approximately equal mass (~1 amu) while electrons are much lighter (~1/1836 amu)
• The number of protons in an atom's nucleus determines its atomic number and element identity
• Atoms of the same element with different numbers of neutrons are called isotopes
• Electrons occupy discrete energy levels or shells around the nucleus with each shell having a maximum number of electrons
  • The first shell can hold up to 2 electrons, the second shell up to 8 electrons, the third shell up to 18 electrons, and so on
• Valence electrons are those in the outermost shell and are involved in chemical bonding

Periodic Table Organization

• Elements are arranged in order of increasing atomic number (number of protons) from left to right and top to bottom
• Vertical columns are called groups or families and contain elements with similar chemical properties due to having the same number of valence electrons
• Horizontal rows are called periods and represent the number of electron shells an element has
• The periodic table is divided into four blocks based on the subshell in which the last electron is added
  • s-block (groups 1 and 2), p-block (groups 13-18), d-block (groups 3-12), and f-block (lanthanides and actinides)
• Metals are located on the left side of the periodic table, nonmetals on the right side, and metalloids along the zigzag line separating metals and nonmetals
• Alkali metals (group 1), alkaline earth metals (group 2), halogens (group 17), and noble gases (group 18) are important element families with distinct properties
• The periodic table helps predict an element's chemical behavior, bonding tendencies, and physical properties based on its location

Electron Configuration

• Electron configuration describes the arrangement of electrons in an atom's shells and subshells
• Electrons fill orbitals in order of increasing energy following the Aufbau principle, Hund's rule, and the Pauli exclusion principle
  • Aufbau principle states that electrons fill orbitals from lowest to highest energy (1s, 2s, 2p, 3s, 3p, 4s, 3d, etc.)
  • Hund's rule states that electrons fill orbitals of equal energy singly before pairing up with opposite spins
  • Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms)
• Electron configurations are written using the element's symbol followed by the number of electrons in each subshell (e.g., Na: 1s²2s²2p⁶3s¹)
• Noble gas notation abbreviates electron configurations by replacing the core electrons with the nearest noble gas symbol (e.g., Na: [Ne] 3s¹)
• Valence electrons are represented by the electrons in the outermost shell and determine an element's chemical properties and bonding behavior
• Elements in the same group have similar electron configurations and chemical properties due to having the same number of valence electrons

Types of Chemical Bonds

• Chemical bonds are attractive forces that hold atoms together in compounds and molecules
• Ionic bonds form between a metal and a nonmetal through the transfer of electrons
  • The metal loses electrons to form a positively charged cation, while the nonmetal gains electrons to form a negatively charged anion
  • The resulting oppositely charged ions are held together by strong electrostatic attractions (e.g., NaCl, MgO)
• Covalent bonds form between two nonmetals through the sharing of electrons
  • Atoms share one or more pairs of electrons to achieve a stable octet configuration
  • Single, double, and triple bonds involve the sharing of one, two, or three electron pairs, respectively (e.g., H₂, O₂, N₂)
  • Polar covalent bonds occur when electrons are shared unequally due to differences in electronegativity (e.g., H₂O, HCl)
• Metallic bonds occur between metal atoms where valence electrons are delocalized and shared among all atoms
  • The delocalized electrons form an "electron sea" that holds the metal cations together
  • Metallic bonding explains the properties of metals such as high electrical and thermal conductivity, malleability, and ductility
• Hydrogen bonds are a type of intermolecular force that occurs between a hydrogen atom bonded to a highly electronegative atom (N, O, or F) and another electronegative atom
  • Hydrogen bonds are stronger than other intermolecular forces and contribute to the unique properties of water and the structure of DNA and proteins

Periodic Trends

• Periodic trends are patterns in the properties of elements that occur across a period or down a group in the periodic table
• Atomic radius decreases from left to right across a period and increases down a group
  • Across a period, the increasing number of protons pulls electrons closer to the nucleus, reducing the atomic radius
  • Down a group, the addition of electron shells increases the distance between the nucleus and the outermost electrons
• Ionization energy increases from left to right across a period and decreases down a group
  • Across a period, the increasing effective nuclear charge makes it harder to remove an electron
  • Down a group, the increasing atomic radius and shielding effect of inner electrons make it easier to remove an electron
• Electronegativity increases from left to right across a period and decreases down a group
  • Across a period, the increasing effective nuclear charge attracts electrons more strongly
  • Down a group, the increasing atomic radius and shielding effect of inner electrons decrease the attraction for electrons
• Metallic character decreases from left to right across a period and increases down a group
  • Across a period, the increasing ionization energy and electronegativity make elements less likely to lose electrons and form cations
  • Down a group, the decreasing ionization energy and increasing atomic radius make elements more likely to lose electrons and form cations

Molecular Geometry

• Molecular geometry describes the three-dimensional arrangement of atoms in a molecule
• VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on the number of electron domains (bonding and lone pairs) around the central atom
  • Electron domains repel each other and arrange themselves to minimize repulsion, resulting in specific geometries
  • Lone pairs occupy more space than bonding pairs, causing slight distortions in the geometry
• Linear geometry occurs when there are two electron domains (e.g., CO₂, BeH₂)
• Trigonal planar geometry occurs when there are three electron domains (e.g., BF₃, SO₃)
• Tetrahedral geometry occurs when there are four electron domains (e.g., CH₄, NH₄⁺)
  • Lone pairs on the central atom can distort the geometry to trigonal pyramidal (e.g., NH₃) or bent (e.g., H₂O)
• Trigonal bipyramidal geometry occurs when there are five electron domains (e.g., PCl₅)
  • Lone pairs on the central atom can distort the geometry to seesaw (e.g., SF₄) or T-shaped (e.g., ClF₃)
• Octahedral geometry occurs when there are six electron domains (e.g., SF₆)
  • A lone pair on the central atom can distort the geometry to square pyramidal (e.g., BrF₅)
• Molecular geometry influences the polarity, reactivity, and physical properties of molecules

Real-World Applications

• The periodic table is used to develop new materials with desired properties, such as semiconductors for electronics and catalysts for chemical reactions
• Understanding chemical bonding is essential for designing drugs that interact with specific biological targets, such as enzymes or receptors
• Molecular geometry plays a crucial role in the function of biomolecules, such as the shape of enzymes determining their catalytic activity and the structure of DNA enabling genetic information storage
• Ionic compounds, such as sodium chloride (NaCl) and calcium carbonate (CaCO₃), have applications in food preservation, water treatment, and building materials
• Covalent compounds, such as water (H₂O) and glucose (C₆H₁₂O₆), are essential for life processes, including cellular respiration and photosynthesis
• Metallic materials, such as steel and aluminum, are used in construction, transportation, and manufacturing due to their strength, durability, and conductivity
• Hydrogen bonding is responsible for the unique properties of water, such as its high boiling point, surface tension, and ability to dissolve polar substances, which are crucial for life on Earth
• Electronegativity differences are used to predict the polarity of molecules, which influences their solubility, melting point, and boiling point, important for separating and purifying compounds