Bonding • Please read and answer the sheet you have been given.
Energy and Chemical Bonds • Chemical bonds are the forces that hold atoms together in a compound. • Energy is required to overcome these attractive forces and separate the atoms in a compound. Thus, the breaking of a chemical bond is an endothermic process. • If energy is required to break a bond, then the opposite process of forming a bond must release energy. The formation of a bond is an exothermic process.
Question • When a chemical bond is broken, the resulting compound has more potential energy than the substance from which it was formed. Why? • Breaking bonds requires energy, therefore the new compound will have more energy than at the beginning. Endothermic AB + energy → A + B
Question • Conversely, when a chemical bond is formed, the resulting compound has less potential energy than the substances from which it was formed. Why? • Forming bonds releases energy, therefore the new compound will have less energy than at the beginning. Exothermic A + B → AB + energy
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Endothermic or Exothermic? Exothermic
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1.) What is a Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • decrease potential energy (PE) • increase stability
CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2
COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3
ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-
Chemical bonds are formed when valence electrons are: • transferred from one atom to another (ionic) • shared between atoms (covalent) • mobile within a metal (metallic)
Sodium has one valence electron and chlorine has seven. Sodium want to lose 1 electron and chlorine needs to gain 1. Sodium transfers its valence electron to chlorine Forming an Na+ and a Cl- ion – sodium chloride NaCl Ionic bonds are formed when metals transfer their valence electrons to nonmetals.The oppositely charged ions attract each other to form an ionic bond.
Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions. atom ion molecular compound ionic compound
Dots represent valence electrons.Everything else (inner shell electrons and nucleus) is called the Kernel and is represented by the symbol. Phosphorous has 5 valence electrons so we draw 5 dots around the symbol for phosphorous.
When metals lose electrons to form ions, they lose all their valence electrons. The Lewis Dot Structure of a metal ion has no dots. The charge indicates how many electrons were lost. Magnesium atom Magnesium ion
When nonmetals gain electrons, they fill up their valence shell with a complete octet (except hydrogen.) The ion is placed in brackets with the charge outside the brackets.
A + metal ion is attracted to a – nonmetal ion (opposites attract) forming an ionic compound. We can use Lewis dot structures to represent ionic compounds. The formula for magnesium fluoride is MgF2
Two major categories of compounds are ionic and molecular (covalent) compounds. • Ionic compounds are formed when a metal combines with a nonmetal. • Ionic compounds have ionic bonds. • Molecular compounds are formed between two nonmetals. • Molecular compounds have covalent bonds.
Comparing the properties compounds with ionic bonds and compounds with covalent bonds. • Properties of ionic compounds • Solids with high melting and boiling points (strong attraction between ions) • Electrolytes: Do not conduct electricity as solids but do when dissolved or molten – ions are charged particles that are free to move • No individual molecules • Properties of molecular compounds • Low melting and boiling points (weak attraction between molecules) • Nonelectrolytes: Do not conduct electricity as solids or when dissolved or molten – no charged particles (ions) to move • Solids are soft • Forms molecules
Ionic solids conduct electricity when dissolved or molten. Molecular solids do not. Solution doesn’t conduct electricity Solution conducts electricity Ionic Solid dissolved in water Molecular Solid dissolved in water
Nomenclature “Or How Do We Name Compounds”
Systematic Naming • Compound is made up of two or more elements • Name should tell us how many and what type of atoms • Too many compounds to remember all the names
Anion • Negative ion • Has gained electrons • Non metals form anions Cation • Positive ion • Formed by losing electrons • Metals form cations
Ionic Compounds • Made of cations and anions • Metals and nonmetals • Electrons lost by the cation are gained by the anion
Na + Cl Ionic Compounds Sodium is cation 1- 1+ Na + Cl Chlorine is anion
Naming Ions • Metal ion is written first in both name and formula • It is named directly from element which formed the ion. • Will nearly always be the positive ion or “cation” • Transition metals can have more than one type of charge • Indicate the charge with roman numerals in parenthesis. Iron(II) or Iron(III) • Exceptions: • Silver always +1 • Cadmium and Zinc always +2
Name these • Na 1+ • Ca 2+ • Al 3+ • Fe 3+ • Fe 2+ • Pb 2+ • Li 1+ • Sodium • Calcium • Aluminum • Iron (III) • Iron (II) • Lead (II) • Lithium
Write Formulas for these • Potassium ion • Magnesium ion • Copper (II) ion • Chromium (VI) ion • Barium ion • Mercury (II) ion • K1+ • Mg2+ • Cu2+ • Cr6+ • Ba2+ • Hg2+
Naming Anions • Anions are always the same. • Change the element ending to -- ide • F1- Fluorine to Fluoride
Name These • Cl1- • N3- • Br 1- • O2- • I1- • Sr2+ • Chloride • Nitride • Bromide • Oxide • Iodide • Strontium
Write These • Sulfide ion • Iodide ion • Phosphide ion • Strontium ion • S2- • I1- • P3- • Sr2+
Polyatomic Ions • Tightly bound groups of atoms acting as a single ion. • Names given in table in book. (pg 123) • Most are anions that contain oxygen. Names end in –ate(one more O), or –ite (one less O). • SO32- = sulfite; SO42- = sulfate • Exceptions: Ammonium cation NH4+, Cyanide CN-, and hydroxide OH-
Naming Binary Ionic Compounds • 2 elements involved • Ionic – metal (cation) and a non-metal (anion) • Naming is easy with representative elements in A groups • NaCl = Na+ Cl-= sodium chloride • MgBr2 = Mg2+Br-= magnesium bromide
Naming Binary Ionic Compounds • The problem comes with the transition metals. • Need to figure out their charges • All ionic compounds will have a neutral charge • Same number of + and – charges • Use the anion to determine the charge on the positive ion.
Naming Binary Ionic Compounds • Try naming these • KCl • Na3N • CrN • ScP • PbO • PbO2 • Na2Se • Potassium chloride • Sodium nitride • Chromium (III) nitride • Scandium (III) phosphide • Lead (II) oxide • Lead (IV) oxide • Sodium selenide
Tertiary Ionic Compounds • Will have polyatomic ions • At least 3 elements • Use blue sheet • Name these ions • NaNO3 • CaSO4 • CuSO3 • (NH4)2O • LiCN • Fe(OH)3 • (NH4)2CO3 • NiPO4 • Sodium nitrate • Calcium sulfate • Copper (II) sulfite • Ammonium oxide • Lithium cyanide • Iron (III) hydroxide • Ammonium carbonate • Nickel (III) phosphate
Polyatomic ions are groups of atoms covalently bonded together that have a negative or positive charge.
Polyatomic ions are held together by covalent bonds but form ionic bonds with other ions. + Ionic bond H - Covalent bonds H N H Cl H
Writing Formulas • Charges have to add up to zero. • Get charges on pieces from Periodic Table • Cations from element name on table • Anions from table change ending to –ide, or use name of polyatomic ion • Balance the charges • Put polyatomics in parenthesis
Writing Formulas • Write formula for calcium chloride • Calcium is Ca2+ • Chloride is Cl1- • Ca+2Cl-1would have a +1 charge • Need another Cl1- • Ca+2Cl2-1 = CaCl2
Writing Formulas • Crisscross method Calcium chloride CaCl2 Ca2+ Cl1- No need to write the one Iron (III) sulfide Fe 3+ S2- Fe 2 S3 Fe2S3
Write Formulas for These • Lithium sulfide • Tin (II) oxide • Tin (IV) oxide • Magnesium fluoride • Copper (II) sulfate • Iron (III) phosphide • Iron (III) sulfide • Ammonium chloride • Ammonium sulfide • Li2S • SnO • SnO2 • MgF2 • CuSO4 • FeP • Fe2S3 • (NH4)Cl • (NH4)2S