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Unit 10 – Chemical Bonding & Nomenclature
At an atomic level, the properties and form of each atom affects the types of relationships and interactions with other atoms and compounds. 1. What are the ways in which types of bonding (ionic, covalent, metallic) differ form one another? What are specific properties of each? 2. How is electron behavior related to each type of bonding? How does it affect the properties? 3. Which elements form which types of bonding? 4. What are the rules of nomenclature for ionic and covalent bonding?
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Chemical Bonds Chemical Bond - an attractive force that holds atoms together in a ___________ Three types of bonding can occur Ionic Bonding Covalent Bonding Metallic Bonding compound
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Ionic Bonding negative
Ionic Bond - electrostatic _________ between ___________ and positive ions Ions - ___________ atoms Atoms that have either __________or lost ___________ electrons (outer level electrons) attraction Charged gained valence
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Ionic Bonding lost Positive Ions - atoms that have _______ electrons (now have _______ protons than electrons) metal atoms Negative Ions - atoms that have ________ electrons (now have _______ electrons than protons nonmetal atoms more gained more
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Energy and electrons in ionic bonding
Reaction energy released = heat of formation Divided conceptually into half-reactions Electron transfer rules for single atoms Electrons lost/gained to form ___________ octets Number of electrons gained is equal to the number of electrons __________ complete lost
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Ionic Bonds Chemical bond due to ______________ __________________
Form crystalline solids with __________ geometric structure Example: NaCl – Sodium Chloride Na ______________; Cl ____________ No single NaCl molecule - ______________ electrostatic attraction orderly loses an electron gains an electron formula unit
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Ionic Compounds Ionic compounds Characterized by ______ bonds
White, crystalline solids soluble in water Families IA and IIA lose electrons and form __________ ions Families VIA and VIIA gain electrons to form __________ ions ionic positive negative
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Chemical Formulas Gives the elements in compound and their proportions for a _____________ _______ Proportions determined by the ____________ of the ions that form the compound For ionic compounds, the net charge of a formula unit is _______ formula unit charges zero
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Ionic Compound Formulas
Two rules Write symbol for __________ ion first followed by __________ ion symbol Assign _________ to assure compound is electrically __________ Example: Calcium chloride positive negative subscripts neutral
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Writing Formulas for Ionic Compounds
Write the formulas for the following ionic compounds: Potassium iodide ___________ Magnesium fluoride __________ Aluminum bromide _________ Lithium sulfide ___________ KI MgF2 AlBr3 Li2S
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Ionic compound names - Nomenclature
metal Name of _____________ (positive) ion first; then __________ (negative) ion, usually with an _______ ending added to the name Ex. oxygen ________ Some transition metals have ____________ charges Historical suffix usage “-ic” for higher of two; “-ous” for lower Modern approach – English name of metal followed by Roman numeral indicating charge nonmetal -ide oxide various
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Examples Write the formulas for the following compounds:
Zinc chloride __________ Iron (III) chloride _________ Iron (II) oxide _________ Give the names of the following compounds CuBr2 ____________________________ FeO ________________________________ CuCl ________________________________ ZnCl2 FeCl3 FeO Copper (II) bromide Iron (II) oxide Copper (I) chloride
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Polyatomic Ions Some ions are formed by groups of atoms that are bonded together These groups are referred to as ___________________ __________ “poly-” means ______________ These polyatomic ions carry a __________ and form ionic bonds just like single ions The names and charges of the polyatomic ions do not need to be memorized and are used in naming the compounds and determining the formula of the compound. When 2 or more polyatomic ions present, must be in parentheses. polyatomic ions many charge
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Examples Write the formulas for the following compounds:
Potassium nitrate ____________ Magnesium sulfate ______________ Sodium carbonate _____________ Lead (II) nitrate ______________ Give the names of the following compounds KOH ________________________________ Na2SO4 _____________________________ FePO4 ______________________________ Fe3(PO4)2 ____________________________ KNO3 MgSO4 Na2CO3 Pb(NO3)2 Potassium hydroxide Sodium sulfate Iron (III) phosphate Iron (II) phosphate
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Covalent Bonds pairs Chemical bonds formed by ______ of ______________ electrons Electrons shared to form ________, ideally ________________ of shared electron clouds between nuclei yields net attraction Atoms within covalent compounds are electrically _____________, or nearly so Octets achieved through ___________electrons Typically between ___________ elements valence octets Overlapping neutral shared nonmetal
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Compounds and chemical change
Atom - smallest ____________ unit Molecule – smallest particle still retaining the characteristic ___________ properties of a substance Examples: Oxygen or hydrogen gas - _____________ molecules _________ - triatomic oxygen molecule Noble gases exist as _____________ particles elemental chemical diatomic Ozone monatomic
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Covalent compounds and formulas
Covalent compound - held together by ______________ bonds Electrons __________ in covalent bonds Electron dot representation Bonding pairs shared _________ (non-bonding pairs) not shared covalent shared Lone pairs
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Multiple bonds Sharing of more than one electron pair Examples double
Ethylene - ____________ bond Acetylene - ______________ bond double triple
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Multiple bonds
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Covalent compound names - Nomenclature
2 Composed of ___ or more nonmetals Same elements can combine to form a variety of different compounds Examples: carbon dioxide or carbon monoxide Two rules First element in formula named with number of atoms indicated by ________ prefix Second element uses Greek prefix for number of atoms of that element and name of element ending in _________ Never use “mono” on first element if only 1 Greek -ide
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Covalent compound formulas
Valence electrons needed determines the number of covalent bonds an atom can form Hydrogen: valence electrons = _____ Forms ___________ bonds only Oxygen: valence electrons = _____ Forms __________ and _________ bonds Nitrogen: valence electrons = _____ Forms _________, __________ and __________ bonds Carbon: valence electrons = _____ No ______________ bonds 1 single 2 single double 3 single double triple 4 single double triple quadruple
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Name the two compounds shown above
______________________________ Carbon dioxide Carbon tetrachloride
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Give the names of the following compounds
NO2 – _____________________________ C2H6 – ____________________________ N2O4 – ____________________________ SO2 – _____________________________ Determine the formulas of the following compounds Carbon tetrabromide _________________ Dinitrogen tetrahydride _________________ Tetraphosphorus pentoxide _______________ Tricarbon octahydride ________________ Nitrogen dioxide Dicarbon hexahydride Dinitrogen tetroxide Sulfur dioxide CBr4 N2H4 P4O5 C3H8
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Lewis Structures Lewis structures are representations of molecules showing all _____________ electrons, bonding and _______________. valence non-bonding pairs
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Writing Lewis Structures
total Find the ________ of valence electrons of all atoms in the molecule. Use the ________________ to determine the number of valence electrons Periodic table
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Writing Lewis Structures
PCl3 (7) = 26
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Writing Lewis Structures
least The central atom is the ___________ electronegative element that isn’t _____________. Connect the outer atoms to it by __________ bonds, using a straight line. Each line represents the two electrons that are shared (-6). hydrogen single
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Writing Lewis Structures
Keep track of the electrons: 26 6 = 20
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Writing Lewis Structures
outer Fill the octets of the _______ atoms. (+18)
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Keep track of the electrons:
26 6 = 20 18 = 2
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Writing Lewis Structures
central Fill the octet of the _______ atom. (+2)
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Keep track of the electrons:
26 6 = 20 18 = 2 2 = 0
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Writing Lewis Structures
If you run out of electrons before the central atom has an octet, form ___________ bonds (double or triple bonds only) until it does. multiple
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Examples Cl2 Chlorine ____________ Cl Cl
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O O C Examples CO2 Carbon Dioxide ____________ (2 x 6) + 4 =
16 valence electrons
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O C O Examples CO2 Carbon Dioxide ____________
Does the central atom have an octet yet? Now what? 4 electrons in the two bonds: 16 – 4 = 12 12 – 12 = 0
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Examples Carbon Tetrachloride ____________
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Examples Nitrogen Trihydride _____________
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Examples C2H4 – ____________________
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Bond polarity unequal Result of ___________ sharing of electrons
Electronegativity – Measure of an atom’s ability to ___________ electrons Differences: 1.7 or greater - ___________ ___________________ Less than ________________ attract ionic polar covalent covalent
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