Chemical Bonds Review PDF

The Ionic Model of Bonding Formation of Ions Octet Rule: - During bond formation, atoms tend to gain a valence shell with a total of 8e- - Metals: lose e- to form cations - Nonmetals: gain e- to form anions Transition Metal Ions: - Can form more than one ion Formati...

The Ionic Model of Bonding Formation of Ions Octet Rule: - During bond formation, atoms tend to gain a valence shell with a total of 8e- - Metals: lose e- to form cations - Nonmetals: gain e- to form anions Transition Metal Ions: - Can form more than one ion Formation Ionic Compounds - E-s are transferred from one atom (usually a metal atom) to another (usually a nonmetal atom) to form a cation and anion - Ionic bonding forms due to the electrostatic attraction between oppositely charged ions (cation and anion) Ionic Compound Nomenclature: - Includes naming (ex. Sodium chloride) and writing formula (ex. NaCl) A. Binary Compounds Steps for Naming 1. Name the metal ion first 2. Name the nonmetal ion (change the suffix to “ide”) Ex. BaS → barium sulfide Steps for Writing Formula: 1. Check if you have a metal and a nonmetal on the periodic table 2. identify the the atomic symbol and ion changes 3. Balance charges 4. Write the formula (metal and then nonmetal) Ex. aluminum oxide → Al2O3 B. Polyatomic ions Nomenclature: Polyatomic ions: Ammonium NH4+ Hydroxide OH− Nitrate NO3– Nitrite NO2– Ex. iron (II) nitrate → Fe(NO3)2 Hydrogencarbonate HCO3– Carbonate CO3^2− Carbonite CO2^2− Sulfate SO4^2– Ex. cobalt (III) sulfate → Co2(SO4)3 Sulfite SO3^2– Phosphate PO4^3– Phosphite PO3^3– Chlorite ClO2– Chlorate ClO3– C. Acid Nomenclature Binary acid - Hydro__ic acid ie. HCl → hydrochloric acid Oxyacid - __ite ion __ous ion ie. HClO2 → chlorous acid - __ate acid __ic acid ie. HClO3 → chloric acid Memory tool: I ate something icky Dynamite is dangerous D. Hydrates Nomenclature - When a crystal of an ionic compound is grown by evaporation from aqueous solution, frequently the crystal structure will include water molecules Ex. CuSO4 · H2O = copper (II) sulfate monohydrate Iron (III) bromide hexahydrate = FeBr3 · 6H2O 1. Mono 2. Di 3. Tri 4. Tetra 5. Penta 6. Hexa 7. Hepta 8. Octa 9. Nona 10. Deca Structure of Ionic Compounds - An ionic crystal consists of three dimensional lattice of opposite ions - Lattice = repeating pattern of the empirical formula - Ionic compounds have giant structures - No individuals molecules - all positive ions attract all the negative ions - Ionic bonding is electrostatic attractions between oppositely charged ions Physical Properties of Ionic Compounds Melting point/boiling point - Tend to have high melting and boiling points - Because the electrostatic forces throughout the giant lattice must be (partially) overcome, and since the electrostatic forces between oppositely charged ions are strong, a lot of energy is required Volatility: the tendency of a substance to vaporize - Tend to have low volatility - Because the electrostatic forces between ions are strong, therefore, requires a lot of energy to overcome Solubility in water - Often soluble in water - When dissolved, individual ions are present, so the lattice structure has to break apart - A lot of energy is required because of the strong forces of attraction between oppositely charged ions in the lattice  - Water is polar, and energy is released when the ions are hydrated by surrounding water molecules (ion-dipole attraction) This energy pays back the energy required to break apart the ionic lattice Solubility in non-polar solvents - Not usually soluble in non-polar solvents - Because only weak forces would be formed between ions and non-polar solvents, this would not pay back the large amount of energy required to break apart the ionic lattice Electrical Conductivity - Recall: electricity is the flow of charged particles (either e- or ions) - As solid, don’t conduct E. because ions are held tightly in position in the lattice structure and are not free to move - As liquids (molten), conduct E. because the forces between the ions are partially overcome and can move freely - When dissolved, conduct E. because the ion are separated from each other and free to move (cations move towards negative electrode, anions to positive electrode) Lattice Enthalpy and Strength of Ionic Bonding - Enthalpy = energy - Lattice enthalpy = the energy required to break apart 1 mol of an ionic solid into its constituent gaseous ions - NaCl (s) → Na+ (g) + Cl– (g) - Energy has to be supplied to break apart the lattice because of the electrostatic attraction between oppositely charged ions - High lattice enthalpy = high melting point - Trend: smaller ions → increasing charge density (ion charge/volume) → stronger ionic bonding → higher lattice enthalpy → higher MP and BP - Trend: increased magnitude of charge on ions → increased electrostatic attraction → stronger ionic bonding → increased lattice energy → higher MP and BP The Metallic Model of Bonding Metallic Bonding: the electrostatic attraction between positive ions and delocalized e- - Metals contain a regular lattice structure arrangement of positive ions (cations) surrounded by a sea of delocalized e- - Delocalized e- don't belong to any one metal but move between all the metal ions in the lattice Metal Properties and Uses Electrical Conductivity - Good conductors of electricity in the solid and liquid state - Delocalized e- are free to move Thermal Conductivity - Good conductors of heat - Because delocalized e- gain kinetic energy where heat is applied and begin moving faster - Move throughout the lattice and transfer their kinetic energy quickly to other electrons and metal ions in the lattice Malleability/Ductility - Malleable and ductile  - Metallic bonding is non-directional (the metals ions in the lattice attract delocalized e- in all directions) - Force applied → layers of ions slide over another without affecting bonding - Alloys: made by dissolving one or more metals in another molten (liquid) metal - Ex steel (iron + carbon), brass (copper + zinc), bronze (copper + tin) - Alloying a metal usually enhances its properties - Stronger and more chemically stable than pure metals - Because mixing different-sized atoms disrupts the lattice structure and prevents planes of metal atoms from sliding over each other as easily. - Less malleable than their constituent metals Strength of Metallic Bonding - Depends on the number of delocalized electrons, charge/radius of the cation - Impacts the melting point Trends in MP down a group - Decreases - Radius get larger down the group, the distance between nucleus and delocalized electrons increases → electrostatic attraction decreases → less energy is required to break apart the lattice structure