Nomenclature and Bonding

Ionic vs Covalent

Remember: the octet rule is filled by either sharing or transferring electrons ○ Based on electronegativity differences between the elements involved ● > 1.7 – electrons are transferred ● ≤ 1.7 – electrons are shared

> 1.7 – electrons are transferred → Ionic Bond ● ≤ 1.7 – electrons are shared → Covalent Bond

Diatomic Elements

diatomic elements: the atoms appear in pairs 1. Hydrogen (H) 2. Nitrogen (N) 3. Oxygen (O) 4. Fluorine (F) 5. Chlorine (Cl) 6. Bromine (Br) 7. Iodine (I)

Molecular Formula ● Molecule: two or more atoms that have been chemically combined ● Compound: substance that contains two or more elements chemically combined in a fixed proportion ● Molecules are small, at the atomic level ● Compounds are visible to the naked eyearrangement of elements depend on particular structure ● number of each atom is indicated by subscript following the atom (if only one atom no subscript written) ● does not tell us about shape of molecule or location of the different atoms ● C12H22O11Chemical Formula: symbols showing elements present in a compound and their relative proportions ● Molecular Formula: chemical formula of a molecular compound showing type and number of atoms in a molecule of the compound ● NH3 H2O C6H12O6 KCl

Empirica Formula

Empirical Formula: shows elements in a compound in their lowest whole-number ratio ● C6H12O6 — CH2O ● many compounds have molecular and empirical formulas that are the same ● H2O P2O5

determination: ○ find percentage of each element ○ find grams of each in 100 g of original material ○ number of moles of each element ○ ratio reduced to small whole numbers ● for molecular, determine molecular weight of compound

Binary Ionic Compounds

compound composed of a monatomic metal cation and a monatomic nonmetal anion ● metal cation first element in formula, nonmetal anion second ● subscripts do not affect the name ● overall charge compound is zero ● metal: element name ● non-metal: drop the end, add -ide

Criss Cross Method 1. write symbol for each ion, including charge 2. cross charges to subscript, disregarding sign 3. reduce subscripts to lowest terms ● ex: calcium iodide; barium oxide

Stock Name System

naming ionic compounds involving transition metal cations ● use roman numerals to indicate the charge of cation

Polyatomic Ions

Ion composed of more than one atom ○ tightly bonded together so entire ion behaves as single unit ○ covalently bonded set of two or more atoms ○ has a net +/- charge

Insert name of ion where it would go in an ionic compound ● Majority are anions which end in -ate or -ite ○ greater #O atoms: -ate → this the starting base ○ lesser #O atoms: -ite ○ one more O than -ate: per- ○ one less O than -ite: hypo-

Ternary Ionic Compounds

ionic compound composed of three or more elements ● one type of cation and one type of anion ● cation, anion, or both is a polyatomic ion ● naming is same as binary ionic compounds ○ cation named first, then anion

when more than one polyatomic ion present, formula is in parentheses with subscript outside indicating how many ions are in the compound ● NaNO3 ● NH4Cl ● Fe(OH)3

Binary Molecular Compounds

Molecular compounds: take form of discrete molecules joined by covalent bonds ○ no ions form; its atoms sharing their valence electron in such a way that a bond forms between the pair of atoms ○ ie: H2O CO2 ○ in large sample of a molecular compound, all the individual molecules are identical

recall: molecular formula show the number of atoms of each element that a molecule contains ○ water: H2O — two hydrogen and one oxygen ○ octane: C8H18 — eight carbon and eighteen hydrogen

binary molecular compound: composed of two elements ○ both nonmetal atoms ● ionic charges cannot be used to name them or write their formulas; they do not reduce either ● two nonmetal atoms will frequently combine with one another in a variety of ratios ○ ie. NO, NO2 , N2O — can’t all be called nitrogen oxide

Generally, less electronegative element is written first in formula (there are a few exceptions). ○ Carbon is always first in a formula and ○ hydrogen is after nitrogen (ie NH3 ). ○ Order of common nonmetals in binary compound formulas is C, P, N, H, S, I, Br, Cl, O, F 2. Prefix is used only if there are more than one atom of that element in the formula

The second element is named after the first, but with the ending changed to —ide. Prefix always used for second element 4. The a or o at end of prefix is usually dropped from name when element begins with vowel (ie 4 O — tetroixde not tetraoxide)

Chemical Bond Formation

chemical bonds: attractive forces between atoms or ions ● form when atoms share or transfer valence electrons ● energy is released when a bond forms; energy is absorbed when a bond breaks — therefore, atoms in bond are at lower energy

Metallic Bond

force of attraction between a positive metal ion and valence electrons surrounding it — both its own valence electrons and those of the other ions of the same metal ● Form lattice-like structure (crystalline solids) where every point in crystal lattice is occupied by an identical atom

Valence electrons are mobile within the solid ● positive cation particles with delocalized electrons drifting between one atom to another ○ this makes the metallic bond — the attraction of the stationary metal ion with the surrounding mobile valence electrons

strong bond produces rigid solids with a definite shape ● high melting points — lots of energy needed because attractive force is so high ● luster and conductivity of electricity caused by mobile electrons ● malleable because if force is applied, electrons can slip in between cations to avoid them coming next to each other

Ionic Solids

Lattice structure ● Properties: solid at STP, high melting point, brittle, poor conductors (unless fused/melted→ liquid state), soluble in water (polar)

Molecular Compounds

Also called covalent compounds ● Shared valence electrons are attracted to nuclei of both atoms ● different elements —> covalent compound ● same elements —> diatomic elements

many burn easily, especially if containing C & H ● many do not dissolve in water ● do not have freely moving electrons — can’t conduct electricity ● individual molecules are more easily separated than ions in a crystal — relatively low boiling points

Multiple Covalent Bonds

some molecules must make more than one covalent bonds ● double covalent bond: formed by atoms that share two pairs of electrons ○ ie C2H4 ● triple covalent bond: formed by atoms that share three pairs of electrons ○ ie N2

make note of C: usually 1s2 2s2 2p2 but this leads to two covalent bonds and an unstable compound ● It will become 1s2 2s12p3 ● ready to make a covalent bond (up to four single bonds or a combination of single and multiple bonds)

Try giving the structural molecules and dot diagrams for the following: ● Acetylene, C2H2 ● NH3 ● H2CO ● BrCl ● O3 ● N2

Network Solids - Macromolecules

Network molecule (network solid/macromolecule): group of atoms covalently bonded together in a large sample. ● Very hard, very high melting point, poor conductor of heat and electricity ● ie Diamond (C), Silicon (Si), Quartz (SiO2 ) ● remember: a molecule is defined as atoms covalently bonded together — therefore a diamond, no matter how big, can be considered a single molecule

Intermolecular Forces

intermolecular forces: forces or attractions between molecules that hold a sample together ● Van der Waals Forces: forces caused by the interaction of electron clouds between molecules; are the interactions between atoms and molecules that result in a pull between them. 1. London Dispersion 2. Dipole-dipole 3. Hydrogen bonding

London Dispersion Forces ○ all materials have this; weakest type ○ attractions arise from the short-term and induced dipoles ○ only significant in noble gases and nonpolar molecules ○ strength can be increased by getting molecules close to one another ○ larger molecules also have stronger forces because greater number of electrons to interact

Dipole-Dipole Attraction: ○ when molecules in sample, oppositely charged ends (dipoles) are attracted to one another ○ attractions cause materials to have higher boiling and melting points ○ stronger type of Van der Waals

attraction between the hydrogen end of a polar covalent molecule to the extremely electronegative end of a neighboring molecule (O, F, or N) ○ special, even stronger form of dipole-dipole attraction ○ unusually high melting and boiling points ○ ie H2O, HF, and NH3