Unit 7 â€“ Introduction to Bonding and Chemical Formulas
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Unit 7 – Introduction to Bonding and Chemical Formulas Addison Wesley – Ch 6, 15, and 16 Practice Forming Ions Atoms lose or gain electrons to become stable Do dot diagram first to determine what will happen Do what is easiest to get 8 electrons Cation – lithium Anion - oxygen Ionic Bond Electrons are everywhere – static is a good example Positive ion is attracted to a negative ion in an ionic bond What kind of elements? Ionic Compound Made up of ions Electrically neutral Charges must equal each other Properties of Ionic Compounds Table salt is a good example High melting point – ionic bonds are strong! Dissolve in water Solution conducts a current Solid doesn’t conduct, molten does Crystal is Brittle Ionic solids shatter along a plane This is because like charges align as soon as they are hit Monatomic Ions Single atom loses or gains electrons to get a stable configuration (octet) Go over charges across table – write on your periodic table Handout Monatomic Cations Positive 1, 2, or 3 Transitions can vary Use element name Use a Roman numeral with any transition that varies Example Copper(I) is Cu+1 Monatomic Anions Can be negative 1,2, or 3 Never vary Change element name to “ide” ending Example: chloride Polyatomic Ions Two or more atoms that are bonded together and carry a single charge Names are on the handout Most are negative with one positive Usually end in “ate” or “ite” Example: NO3- is nitrate Formulas for Binary Compounds Contain a monatomic cation (metal) and a monatomic anion (nonmetal) Metal is first Charges must add to “0” Use subscripts to get the value to “0” Why is sodium chloride NaCl? Try some Shortcut Place charge above each ion Crisscross Reduce if necessary (empirical formula is smallest ratio of ions) Try some Formulas with Polyatomic Ions Compounds are tertiary if they have 3 or more elements Treat polyatomic as a single unit Put in parentheses if you have to Practice Naming Ionic Compounds Use NaCl as a good example Metal first, then nonmetal Ends in “ide” if binary Use polyatomic name if tertiary Use Roman numeral if necessary How do you know? Suspect every transition metal Try some Covalent Bonding Look around Most of what you see is covalently bonded Definition – formed by a pair of electrons that are shared between atoms. This would occur between nonmetals Molecule Atoms joined together by covalent bonds Substance is molecular if it is made up of molecules Examples: CO2, H2O C6H12O6 Examples Empirical Formula This is the simplest whole number ratio of atoms in a substance (reduced) Ionic formulas are always empirical Molecular formulas are sometimes empirical (H2O) and sometimes not (H2O2) Structural Formulas Different formulas may have the same formula, so structural formulas are often drawn Lewis structure is one example Based upon the Lewis dot diagram of elements Try bromine, magnesium Procedure Draw the Lewis dot structure for F2 First, show the dot diagram for each atom involved in the molecule Share electrons in order for each atom to have an octet around it (8 electrons) Try HCl Try single bonds first! Try O2 Double up if you come up with a shortage of one electron. Triple bonds Try N2 Triple up if you end up with a shortage of 2 electrons. Practice: CO2 H 2O H2CO HCN Exceptions Some elements are satisfied with fewer than 8 electrons (6 or 4) Some structures can only be drawn with 7 electrons These substances can be short-lived and reactive – called free radicals Topic of college chemistry Properties of Molecular Substances •Let’s look at these substances •Can be solids, liquids, or gases •Much weaker bonds, lower melting points •Nonconductors •Soft Molecular Shape Try CH4 Build with toothpicks and styrofoam balls Not 2-D! Electrons will repel each other VSEPR Theory Valence Shell Electron Pair Repulsion Theory In a small molecule, the electrons are arranged as far apart as possible Explains most molecules 5 basic shapes Types CH4 – Tetrahedral NH3 – Trigonal Pyramidal H2O – Bent HCl – Linear Mixed Polarity Electrons are often not equally shared Polar bond – Covalent bond between atoms that pull unequally on the electrons How do you know? Electronegativity difference (handout) Example – H-O Polar Try a few Polar Molecule Molecules contains polar bonds Arrangement is not symmetric Try H2O Try CCl4 Some Implications Shape gives molecules its properties (smell) Polarity determines solubility (Milk kaliedoscope) Isomers Molecules with the same formula but different arrangements Three types: Structural (C and H) Functional (alcohol and ether) Stereoisomer (mirror image) Stereoisomers Drug interactions Genetic mistakes Smells (Caraway and Spearmint) Naming Compunds are binary so they end in “ide” Element name first, “ide” form second Prefixes tell how many of each N2O3 - Dinitrogen trioxide Don’t use “mono” if one comes first Try CO2 and CO Write these on your periodic tables: Mono,di,tri,tetra,penta,hexa,hepta,octa, nona,deca Try these! H2O Na2O CuCl N2O3 AlN Check for a metal first!!