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Ionic Nomenclature Recap
____________ is ALWAYS written first. If monoatomic, use the name of the element. If polyatomic, use the name of the polyatomic ion. ____________ is ALWAYS written second. If monoatomic, use -ide as the suffix. Examples: AgCl2 NaOH NH4NO3
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Recall What 2 things are needed to create an IONIC bond: Therefore, what do you think creates a COVALENT bond:
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Covalent Bonding
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Covalent Bonds What happens to electrons in covalent bond?
Electrons ________ between atoms, not given or taken completely. Atomic orbitals are combined to form molecular orbitals. Generally bond between ______________ and _____________. Why? 3 trends need to be taken into account. Electron affinity, Ionization energy, and electronegativity differences between atoms are small. Differences in trends aren’t large enough for one atom to completely take away electrons from the other, so electrons are shared. Example: CH4
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Identification of Bond Types Examples
1.) Na—Cl 11.) Al—Cl 2.) Mg—Cl 12.) N—O 3.) C—N 13.) Be—F 4.) O—Cl 14.) S—C 5.) Zn—Br 15.) Zn—N 6.) Ti—F 16.) I—I 7.) P—C 17.) P—O 8.) K—O 18.) Li—N 9.) F—C 19) C—Cl 10.) Li—I 20.) Ag—I
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Covalent Compounds—Nomenclature
Rules—Binary 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. Do not simplify subscripts in covalent compounds Example— P4O10 Prefixes: Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10 Prefix indicating number of atoms contributed by second element Prefix needed if first element contributes more than one atom Root name of second element + ide Name of first element
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Covalent Compounds—Nomenclature
As2O5 Carbon Tetrafluoride CO Sulfur Trioxide
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BELLRINGER (5 Minutes) SiI4 NO OF2 tetraphosphorus decoxide
dinitrogen tetroxide
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Acids 2 types: Binary Acids—contain H and another element, usually a halogen. 1.) Put hydro- as prefix for H. 2.) Use –ic as suffix for second element. Example—HCl = hydrochloric acid Oxyacids—contain 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 Example—HNO3 = Nitric Acid
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Practice in your notes…
H2SO3 H3PO4 Hypochlorous Acid Arsenic Acid
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Whiteboarding SiI4 NO HF OF2 Carbonic acid hydrosulfuric acid
tetraphosphorus decoxide dinitrogen tetroxide
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Hydrocarbons Hydrocarbons = the simplest organic compounds
Contain only __________ and _________________ 3 types of straight-chain hydrocarbons Alkanes—completely saturated hydrocarbons (no double or triple bonds) Go by the formula CnH2n+2 End in suffix -ane Alkenes—contain a double bond End in suffix -ene Alkynes—contain a triple bond End in suffix -yne
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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
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Bellringer (5 min): Determine type of bond that will be formed
C – O Al – Cl K – Br F – F Mn – O Ca – Se S – Cl Li – S P – Br C – I
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Formation of Covalent Bond
Nature favors chemical bonding…why? Makes the atoms more _______________. potential energy is _________________ when the atoms are bonded. Lower ___________ = more _____________.
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Characteristics of a Covalent Bond
__________________ —the average distance between two bonded atoms. Depends on type of bond: single, double, or triple. Bond _____________ Energy —The energy required to break 1 mol of a specific chemical bond (always _______________). Positive energy to break a bond (endothermic) Negative energy to make a bond (exothermic) Indicates strength of a bond and is directly related to bond length.
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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.
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Molecule—a neutral group of atoms that are held together by covalent bonds.
Molecular formula—shows the types and numbers of atoms combined in a single molecule of a compound. Example: CH4
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Octet Rule Atoms undergo bonding in order to satisfy the ______________. Octet rule: Chemical compounds tend to form so that each atom has ___ electrons, either by ____________, _______________, or _____________ electrons. Atoms want to become “Noble-gas-like” with filled valence shells. Diatomic molecule: A molecule in which there are only two atoms. F2, Cl2, Br2, I2, H2, O2, N2
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Exceptions to Octet Rule
Hydrogen—forms only one bond to have two valence electrons. Group 13—Has 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.
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Lewis Structures What are they?? What are they used for??
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 where electrons are located Gives relative bond strengths to establish reactivity of molecules.
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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 electrons 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 are accounted for. ***** Add all atoms (except H) first, then fill in molecule with Hydrogens
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Practice with Lewis Structures
Draw the Lewis structures for the following molecules: NF3 BH3 PCl5
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Whiteboarding Practice
Draw the Lewis structures for the following molecules: SiH Pl3 H2S CCl2F2 C2F SF6
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Before we go any further…
Try the Lewis structure for ions: Negative charge: add the electrons Positive charge: remove the electrons Ammonium Peroxide Try the Lewis structure for C2H4.
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Single and Multiple Covalent Bonds
Single bond – two electrons shared Also called sigma bond, or bond. Examples: H2 Double bond – four electrons shared Also called pi bond, or bond. Example: O2 Triple bond – six electrons shared Consists of one and one bond. Example: N2
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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
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Relative Bond Lengths and Strengths
Bond length—the more shared pairs = the ___________ the bond. Single > Double > Triple Bond Strength—the more shared pairs = the __________ the bond. Triple > Double > Single
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Hydrocarbons Hydrocarbons = the simplest organic compounds
Contain only __________ and _________________ 3 types of straight-chain hydrocarbons Alkanes—completely saturated hydrocarbons (no double or triple bonds) Go by the formula CnH2n+2 End in suffix -ane Alkenes—contain a double bond End in suffix -ene Alkynes—contain a triple bond End in suffix -yne
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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
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Lewis Structure Bellringer
H2O NH3 PCl5 BF3 SF6 CH4 CS2 HCl
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Resonance Structures In molecules with multiple bonds, electrons are ___________________ due to shared orbitals. Since electrons are delocalized, a single lewis structure cannot account for all of the possible locations of electrons. Only electron locations vary, NEVER atom arrangement.
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Formal Charges Formal Charges are used when there is more than one possible Lewis structure for a molecule (resonance structures). They help us determine with Lewis structure will most likely exist in the real world. The most stable Lewis structure is one where all charges equal zero or the negative charge resides on the most electronegative element. FORMAL = VALENCE NON-BONDING - BONDING ELECTRONS CHARGE ELECTRONS ELECTRONS Example 1: Methane Example 2: Sulfur dioxide
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Polarity Based on electronegativity differences. _______ molecule = uneven distribution of electrons (non-symmetrical) _________________ = even distribution (symmetrical)
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Polar Covalent Bonds
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Non-polar and Polar Covalent Bonds
Dipole—a molecule that contains both positively and negatively charged regions (_______________ sharing of electrons). Nonpolar bonds do not contain a dipole (similar electronegativies). Least Most electronegative electronegative
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Identify the dipole in the following molecules… and indicate if molecule is polar or non-polar
If geometry is symmetrical (NO lone pairs), dipoles will cancel NH3 CH3F SF6 AlCl3 CCl4
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Identify the bond as polar or nonpolar
Remember: C—H Polar Covalent = O—Cl Nonpolar Covalent = 0-0.4 C—Cl H—N B—F F—F
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