Ionic Nomenclature Recap - PC\|MAC

Ionic Nomenclature Recap - PC\|MAC

Ionic Nomenclature Recap Cation is ALWAYS written first. If monoatomic, use the name of the element. If polyatomic, use the name of the polyatomic ion. Anion is ALWAYS written second. If

monoatomic, use -ide as the suffix. If polyatomic, use the name of the polyatomic ion. Covalent Bonding Covalent Bonds What happens to electrons in covalent bond? Electrons Atomic orbitals are combined to form molecular orbitals.

Generally Why? shared between atoms, not given or taken completely. bond between nonmetal and nonmetal. 3 trends need to be taken into account. Electron affinity, Ionization energy, and electronegativity differences between atoms are small. Differences in trends arent large enough for one atom to completely take

away electrons from the other, so electrons are shared. Example: CH4 Moleculea neutral group of atoms that are held together by covalent bonds. Molecular formulashows the types and numbers of atoms combined in a single molecule of a compound. Example: Concept

CH4 check: What is the difference between a molecule and a polyatomic ion? Formation of Covalent Bond Nature favors chemical bondingwhy? Makes the atoms more stable. potential energy is lowered when the atoms are

bonded. Lower energy = more stable. Characteristics of a Covalent Bond Bond lengththe average distance between two bonded atoms. Depends on type of bond: single, double, or triple. Bond Dissociation EnergyThe energy required to break 1 mol of a specific chemical bond (always endothermic). Indicates strength of a bond

In Reactions For endothermic reactions a greater amount of energy is required to break the existing bonds than is released when the new product bonds form. For exothermic reactions more energy is released forming new bonds than is required to break bonds of the reactants. Covalent CompoundsNomenclature RulesBinary

Compounds Element with smaller group # is always given first (similar to cation in ionic bonding). Second element combines prefix with suffix ide. If second element begins with vowel, the o- or a- in the prefix is dropped.

Examplepentoxide Prefixes: Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10 Practice: N2O As2O5 Carbon Tetrafluoride

CO Sulfur Trioxide Acids 2 types: Binary Acidscontain H and another element, usually a halogen. 1.) Put hydro- as prefix for H. 2.) Use ic as suffix for second element. ExampleHCl = hydrochloric acid Oxyacidscontain H, O, and a third element (mostly H paired with a polyatomic ion). 1.) Identify anion. 2.) if the anion suffix is ate replace it with ic 3.) if the anion suffix is ite replace it with ous ExampleHNO3 = Nitric Acid Octet Rule

Atoms undergo bonding in order to satisfy the octet rule. Octet rule: Atoms want to be noble gas-like. Diatomic molecule: A molecule in which there are only two atoms. F2, Cl2, Br2, I2, H2, O2, N2, Exceptions to Octet Rule Hydrogenforms only one bond to have two valence electrons. Group 13Has three valence electrons. Tends to form three bonds.

Some elements can form an expanded octet if bound to highly electronegative atoms. Example: SF6 Expanded octet involves empty d orbitals to fit extra electrons. Lewis Structures What are they?? A formula where atomic symbols represent nuclei and inner shell electrons, and dot pairs represent valence and bonded electrons. What

are they used for?? Gives us a way to visualize bonding between atoms Represents Gives where electrons are located relative bond strengths to establish reactivity of molecules. Lewis Structures Six Steps:

1.) Determine types of atoms in molecule. Example: CH3I 2.) Write electron dot notation for each atom. 3.) Determine the total number of valence electros available. 4.) Arrange atoms with LEAST electronegative atom in the center (exception: H), and place one shared pair of electrons between each of the atoms. 5.) Fill in valence shells of atoms with unshared electrons (lone pairs). 6.) Count electrons to make sure all available valence electrons Practice with Lewis Structures Draw the Lewis structures for the

following molecules: NH3 H2S SiH4 PF3 Before we go any further Try the Lewis structure for C2H4. Single and Multiple Covalent Bonds Single Also bond two electrons shared called sigma bond, or bond.

Examples: Double Consists H2 bond four electrons shared of one sigma and one pi bond. Example: O2 Triple bond six electrons shared Consists

of one and two bonds. Example: N2 Lewis structures with multiple bonds Multiple bonds become evident in lewis structures when there are not enough valence electrons after adding lone pairs. Examples: CH2O

CO2 HCN Hybridization Orbitals of similar energy mix to produce new orbitals (hybrids) Example: Carbon in Carbon tetrafluoride Relative Bond Lengths and Strengths Bond lengththe more shared pairs = the shorter the bond.

Single Bond > Double > Triple Strengththe more shared pairs = the stronger the bond. Triple > Double > Single Resonance Structures In molecules with multiple bonds, electrons are delocalized due to shared

orbitals. electrons are delocalized, a single lewis structure cannot account for all of the possible locations of electrons. Only electron locations vary, NEVER atom arrangement. Polarity Based on electronegativity differences. Polar

molecule = uneven distribution of electrons Nonpolar = even distribution Polar Covalent Bonds Identify the bond as polar or nonpolar FF CH OCl CCl HN BF

Non-polar and Polar Covalent Bonds Dipolea molecule that contains both positively and negatively charged regions (unequal sharing of electrons). Nonpolar bonds do not contain a dipole. Identify the dipole in the following molecules NH3 CH3F SF6 BF3

CCl4 Hydrocarbons Hydrocarbons Contain Can 3 = the simplest organic compounds only carbon and hydrogen be straight-chain, branched chain, or cyclic molecules types of straight-chain hydrocarbons Alkanescompletely triple bonds)

End in suffix -ane saturated hydrocarbons (no double or Need to know first ten alkane hydrocarbons CH4 Methane C6H14 Hexane C2H6

Ethane C7H16 Heptane C3H8 Propane C8H18 Octane C4H10 Butane

C9H20 Nonane C5H12 Pentane C10H22 Decane Properties of Compounds Property Ionic State of Matter

Crystalline solids composed of ions MP and BP Conductivity Hardness High Good in molten or liquid state None in solid state Hard Covalent Solids, liquids, or gases composed of molecules

Low Non-conductors Soft

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