UPCAT Chemistry — Periodic Table, Bonding & Chemical LanguageRevision Notes

Formulas

Exam Tips

• Memorize the charges of common groups: +1, +2, +3, -3, -2, -1 for groups 1,2,13,15,16,17
• Practice identifying element positions by their electron configurations
• Use periodic trends to compare properties without memorizing specific values
• Remember exceptions like hydrogen (can be +1 or -1) and transition metals (variable charges)

Key Points

• The periodic table is arranged in order of increasing atomic number (number of protons)
• Horizontal rows are called periods, vertical columns are called groups or families
• Group 1: Alkali metals (Li, Na, K, Rb, Cs) - all form +1 ions
• Group 2: Alkaline earth metals (Be, Mg, Ca, Sr, Ba) - all form +2 ions
• Group 17: Halogens (F, Cl, Br, I) - all form -1 ions
• Group 18: Noble gases (He, Ne, Ar, Kr, Xe) - very stable, rarely form compounds
• Atomic size decreases across a period (left to right) and increases down a group
• Ionization energy increases across a period and decreases down a group
• Electronegativity increases across a period and decreases down a group
• Metallic character decreases across a period and increases down a group

Definitions

Section Title

The Periodic Table Organization and Trends

Common Mistakes

• Confusing periods with groups - remember periods are horizontal rows
• Mixing up trend directions - atomic size increases down and decreases across
• Forgetting that noble gases don't follow typical ionization patterns
• Not recognizing that transition metals can have multiple oxidation states

Exam Tips

• Use electronegativity differences to predict bond type: >1.7 ionic, <0.4 nonpolar covalent, 0.4-1.7 polar covalent
• Remember IMF strength order: ion-dipole > hydrogen bonding > dipole-dipole > London dispersion
• Practice identifying which IMFs are present in given compounds
• Connect IMF strength to physical properties like boiling point and solubility

Key Points

• Three main types of intramolecular forces: metallic (metal-metal), ionic (metal-nonmetal), covalent (nonmetal-nonmetal)
• Metallic bonds involve a 'sea of electrons' shared among metal atoms
• Ionic bonds form through complete transfer of electrons from metal to nonmetal
• Covalent bonds form through sharing of electrons between nonmetals
• Bond polarity depends on electronegativity difference between atoms
• Intermolecular forces (IMFs) exist between molecules: London dispersion, dipole-dipole, hydrogen bonding, ion-dipole
• London dispersion forces are weakest and exist between all molecules
• Hydrogen bonding occurs when H is bonded to F, N, or O
• Ion-dipole forces are strongest and occur between ions and polar molecules
• Stronger IMFs lead to higher boiling points, lower vapor pressure, higher surface tension, and higher viscosity

Definitions

Section Title

Chemical Bonding Types and Properties

Common Mistakes

• Confusing intramolecular forces (within molecules) with intermolecular forces (between molecules)
• Thinking that only polar molecules have London dispersion forces - ALL molecules have them
• Forgetting that hydrogen bonding requires H bonded to F, N, or O specifically
• Not recognizing that larger molecules have stronger London dispersion forces

Formulas

Exam Tips

• Always start with the least electronegative atom as central (except H)
• Check total electrons: for anions add electrons, for cations subtract electrons
• Use formal charge to evaluate alternative structures
• Practice common geometries: linear (2), bent (2+lone pairs), trigonal planar (3), tetrahedral (4)

Key Points

• Lewis structures show how atoms are connected and where electrons are located
• Steps: count valence electrons, determine central atom, draw bonds, complete octets, assign formal charges
• Central atoms are usually the least electronegative (except hydrogen)
• Some atoms (P, S, Cl and beyond) can expand their octets (hypervalent)
• Formal charge helps choose the best Lewis structure among alternatives
• VSEPR theory predicts molecular shapes based on electron pair repulsion
• Count bonding and nonbonding electron pairs around central atom
• Multiple bonds count as one electron pair region
• Electron geometry considers all electron pairs; molecular geometry considers only atoms
• Lone pairs occupy more space than bonding pairs, affecting bond angles

Definitions

Section Title

Lewis Structures and VSEPR Theory

Common Mistakes

• Forgetting to add/subtract electrons for charged species
• Placing hydrogen as the central atom (it can only form one bond)
• Not completing octets for outer atoms before expanding central atom's octet
• Confusing electron geometry with molecular geometry when lone pairs are present

Exam Tips

• Memorize common polyatomic ions: sulfate (SO₄²⁻), nitrate (NO₃⁻), phosphate (PO₄³⁻), carbonate (CO₃²⁻)
• Learn the criss-cross method for writing ionic formulas quickly
• Practice acid naming patterns with common examples
• Remember that mono- is usually omitted for the first element in covalent compounds

Key Points

• Binary covalent compounds use Greek prefixes (mono-, di-, tri-, tetra-, penta-, etc.)
• First element keeps its name, second element gets -ide ending
• Ionic compounds: name cation first, then anion with -ide ending
• Roman numerals indicate charge for metals with variable oxidation states
• Polyatomic ions have specific names that must be memorized
• Acids: binary acids use hydro-____-ic acid format
• Oxyacids: -ate ions form -ic acids, -ite ions form -ous acids
• Per- prefix for one more oxygen, hypo- prefix for one less oxygen than -ous
• Molecular formula shows actual numbers of atoms
• Empirical formula shows simplest whole number ratio

Definitions

Section Title

Chemical Nomenclature and Formula Writing

Common Mistakes

• Forgetting Greek prefixes for covalent compounds (CO₂ is carbon dioxide, not carbon oxide)
• Using prefixes for ionic compounds (it's calcium chloride, not calcium dichloride)
• Mixing up -ate/-ite and -ic/-ous relationships in acid naming
• Not reducing subscripts to lowest terms when writing formulas