FINAL EXAM REVIEW CHEM 1411

1. FINAL EXAM REVIEWCHEM 1411 You will need to bring a picture ID. Pencils and calculator (nonprogrammable). The Exam is Saturday Dec. 6 at 9:00 am CPC Room TBA The exam is EXACTLY 2 Hours in length. Being late costs you time on the exam. Not taking the final = 0 for Final exam grade THERE ARE NO MAKE-UP FINALS Fall 2008

2. Chapter 1CHEMISTRY: MATTER AND MEASUREMENT • Chemical Changes • Physical Changes • Chemical Properties • Physical Properties • Dimensional Analysis • (Unit Factor Method) • Density

3. Chapter 1CHEMISTRY: MATTER AND MEASUREMENT • Chemical Properties – chemical changes • Changes in chemical composition • Relative reactivity • Physical Properties – physical changes • changes of state, composition remains constant • density, color, solubility • Extensive Properties – depend on amount present • Intensive Properties – physical or chemical properties that are independent of amount present.

4. Chapter 1CHEMISTRY: MATTER AND MEASUREMENT • Density

5. Chapter 2 Atoms, Molecules, and Ions • Isotopes • Atomic Number • Atomic Mass • Chemical compounds • Mixtures • Names and Formulas of Ionic Compounds • Molecular/Formula Weight • Empirical Formulas • Molecular Formulas • Covalent bonds/Covalent compounds

6. Mass Number and Isotopes • The mass number of an element (specifically the isotope of the element) is given the symbol A. • A is sum of protons + neutrons Specific isotopes of an element (E) can be shown as A = mass number Z = atomic number N = neutrons in that isotope

7. Atomic Weights • This is the lower number on periodic chart • This is weighted average of the masses of the constituent isotopes naturally occurring.

8. Mixtures Pure Substances Heterogeneous Homogeneous ChemicalCompounds Elements Compounds and Mixtures Matter

9. Compounds and Mixtures • Mixtures • composed of two or more substances in variable ratios • can be separated into components by simple physical or chemical methods homogeneous mixtures– cannot be separated by physical inspection heterogeneous mixtures– usually can be separated by simple physical means

10. Chemical Bonding • Chemical bonding is the attractive forces that hold atoms together. • Covalent bonds result from the sharing of two (usually) or more electrons between 2 atoms. • Ionic bonds are electrostatic forces describing the attraction between positive and negative ions.

11. Ionic or Covalent? • Is the compound ionic or covalent? • Generally covalent bonds consist of nonmetals (usually) sharing electrons. • Ionic bonds exist between metals and nonmetals (usually) – the metal gave up its electron to the nonmetal. • Polyatomic ions have covalent bonds but also carry a charge.

12. Naming Compounds • By convention, chemical formulas are written with the most positive element listed first. • Binary ionic compounds. For metal ions you only need to indicate the charge on the ion (usually necessary for transition and post-transition metals). The number of that ion present is assumed. • Binary covalent compounds. For nonmetals, a prefix is only used for the cation (first element listed) if more than one atom is present. Since nonmetals may combine in numerous ratios, the anion portion of the name always is given with a numerical prefix.

13. Chapter 3Formulas, Equations, and Moles • Chemical Equations (including balancing) • Calculations Based on Chemical Equations • Limiting Reactant • Concentration of Solutions • Using Solutions in Chemical Reactions • % Composition and emperical formulas • Molecular formulas vs molecular mass

14. with heating reactants or starting materials yields or produces Chemical Equations • Consider a typical reaction: products The coefficients represent the number of moles (or particles) needed for the complete reaction. No coefficient is understood to be “1”. Consider these numbers to be ratios.

15. The Mole • The term mole is a shorthand abbreviation of “gram formula molecular weight” • gram formula molecular weight  gram molecular weight  molecular weight  mol. weight mol.wt.  mol • mol. or mol was spelled mole (like it sounds) • Avogadro’s number = 6.022  1023 You need to make moles your friends (or at least allies) or you will not survive the course.

16. Concentration of Solutions • Nomenclature • solution • a mixture of two or more substances, usually one is a liquid • solvent • the material doing the dissolving (usually the most abundant material or the liquid) • solute • the material that is dissolved (usually the least abundant material) • concentration • the amount of solute dissolved in a solvent, can be expressed in percent (%), molarity (M), normality, (N), or molality (m)

17. Dilution of Solutions • Molarity dilution problems can be easy to solve. • Since the number of moles remains constant M1 = molarity solution 1 V1 = volume solution 1 M2 = molarity solution 2 V2 = volume solution 2 You may use either L or mL (these must be the same units) to solve the problem.

18. Adjust moles Titration Calculations Do the moles change -- NO Do the moles change -- YES

19. Titrations Titration is a method to exactly the determine the concentration of a solution. Titrations require standards – materials whose concentrations can be exactly (accurately) determined special glassware – burets and pipets, these have been calibrated to be accurate to  0.01 mL or less (0.001 mL) indicators – to “indicate” when the titration is finished This allows concentrations to be determined to 4 decimal places (min. 4 significant figures)

20. Percent Composition and Chemical Formulas • Example 24: A compound contains 24.74% K; 34.76% Mn; and 40.50% O by mass. What is its empirical formula? Formula: KMnO4

21. Chapter 4SOME TYPES OF CHEMICAL REACTIONS • Acid/Base Reactions • Precipitation Reactions • Oxidation/Reduction Reactions • Including Oxidation Numbers • Formula Unit Equations (molecular) • Total Ionic Equations • Net Ionic Equations • Balancing Redox reactions

22. Acid-Base Reactions • The reaction of an acid and a base produces a salt and (usually) water. • By definition, neither an acid nor a base (or their anhydrides) can be a salt. Strong acids and bases will dissociate into ions. Weak acids and bases are shown in all representations as the non-dissociated molecules.

23. Strong and Weak Acids • Acids generate H+ aqueous solutions. • The chemical formulas for acids usuallybegin with H. • Organic acids are the exception to this. They end with COOH or CO2H. • Know these formula conventions! • KNOW THE STRONG ACIDS BY MEMORY

24. Strong Soluble Bases • Bases produce hydroxide ions, OH–, in solution. • Strong soluble bases ionize ~100% in water. KNOW THE STRONG BASES BY MEMORY

25. Solubility Rules • For this course, solubility describes a compounds ability to dissolve in water. • You must be able to predict whether a compound is or is not soluble (this is exam material). • All (soluble) strong acids and all strong bases are completely water soluble. • Most but not all of their salts are soluble. • The rules for solubility are used to determine if an inorganic salt is soluble. • You must know these rules by heart!!!

26. Solubility Rules • Soluble Ionic Salts • All IA metal salts are soluble. • All ammonium salts are soluble. • All chlorides are soluble (except Ag, Hg, and Pb). • All nitrates and acetates are soluble. • All perchlorates and chlorates are soluble. • All fluorides are soluble (except IIA metal fluorides). • All sulfates are soluble (except Ca, Ba, Sr, Ag, Hg, and Pb).

27. Solubility Rules • The “solubility” rules should be followed first. • Regardless all IA metal salts and NH4+ are soluble. • Insoluble Ionic Salts • (IA metal and ammonium salts are soluble) • All phosphates are insoluble. • All carbonates are insoluble (exceptMg which is moderately soluble) • All sulfides are insoluble. • All oxides (except the strong acid and base anhydrides).

28. Oxidation Number Rules The oxidation number of any free, uncombined element is zero. The oxidation number of an element in a simple (monatomic) ion is the charge on the ion. In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero. In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion.

29. Oxidation Number Rules Hydrogen, H, in combined form, has the oxidation number +1 (usually). The only exception to this is found with metal hydrides. A compound that is composed of the metal and H. This is the hydride ion, H-. Oxygen, O, in combined form, has the oxidation number –2, (usually). The exceptions to this are found with some metal oxides and hydrogen peroxide. Fluorine as fluoride is always –1.

30. GER Redox Reactions LEO the lion goes “GER.” Loses Electrons – Oxidized Gains Electrons – Reduced

31. Reactions in Aqueous Solutions • There are 3 methods of writing a chemical reaction. • Molecular Equation • This is the typical balanced chemical equation usually seen. It may indicate the state of the reactant: (g) gaseous; (aq) aqueous; (s) solid; (l) liquid – used for water in an aqueous solution. • Total Ionic Equation • This representation shows the reaction as the ions or molecules that are present in the solution. • Net Ionic Equation • This method only shows the ions or molecules that actually change or are directly involved in the chemical reaction.

32. Chapter 5 Periodicity and Atomic Structure • Frequency, wavelength and energy relations • deBroglie wave/particle calculations • Periodic Properties • Electronic Structure of Atoms • The quantum mechanical picture of atoms • Quantum Numbers • Shapes of orbitals • Electron Configuration • Shielding

33. Electromagnetic Radiation • wavelengthl(Greek lower case lambda) • distance from the top (crest) of one wave to the top of the next wave • units of distance - m, cm, Å • 1 Å = 1  10–10 m = 1  10–8 cm • frequency(Greek lower case nu) • this is sometimes represented as  (italicized v) • number of crests (wavelengths) that pass a given point per second • units of frequency = 1/time or s–1 or Hertz (Hz)

34. Electromagnetic Radiation • Relationship for electromagnetic radiation • c = l • c = velocity of light • 3.00  108 m/s or 3  1010 cm/s

35. The Wave Nature of the Electron • Louis de Broglie postulated that electrons have wave-like properties • The wavelengths of electrons are described by the de Broglie relationship.

36. Quantum Numbers • Quantum numbers are solutions of the Schrödinger, Heisenberg & Dirac equations • Four quantum numbers are necessary to describe the energy states of electrons in atoms n –the principle quantum number  – subsidiary quantum number m – magnetic quantum number ms – spin quantum number

37. ElectronicConfigurations • Pauli Exclusion Principle • No two electrons in an atom can have the same set of 4 quantum numbers. • Aufbau Principle • The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available. • Hund’s Rule • Electrons will occupy all orbitals singly before pairing can begin. • The spins of these electrons will be aligned.

38. radii increases radii increases Atomic Radii (A Group Elements Only) • Atomic radii describes the size of atoms. • This increases as you go from the right to the left. (opposite what you would think since more e– are added.) • This increases from top to bottom. • (as expected – more shells are being added)

39. Atomic Radii • The decreasing radii across a period is due to the shielding or screening effect of the inner electrons [He] or [Ne], etc. • Consequently the outer electrons feel a stronger effective nuclear charge than expected. • Li [He] shields effective charge is +1 • Be [He] shields effective charge is +2 • F [He] shields effective charge is +7 • Na [Ne] shields effective charge is +1

40. Chapter 6CHEMICAL PERIODICITY • Periodic Properties of the Elements • Atomic Radii • Ionization Energy • Electron Affinity • Ionic Radii • Electronegativity

41. Chemical Properties • Chemical properties and the reactivities of various elements can be predicted based upon their electronic configurations. • Only the outermost electrons or valence electrons are involved in chemical reactivity and properties.

42. IE decreases IE decreases Ionization Energy - Potential • Ionization energy, IE, generally increases as you go across a period (smaller radii, e– more tightly held). • Important exceptions at Be and Mg, N, and P, because of Hund’s rule Ionization energy generally decreases as you go down a group (larger radii, e– more loosely held).

43. Ionic radii increases Ionic radii increases Ionic Radii • Ionic radii increases as you go from right to left across the periodic table. Ionic radii increases as you go from down the periodic table.

44. Electronegativity increases EN increases Electronegativity • Electronegativity, EN, is the measure of the tendency of an atom to attract electrons to itself in compounds (to gain electrons). • EN increases as you go from left to right (to F). EN increases as you go up the table.

45. Net reaction Born-Haber Cycle for NaCl -348.6 kJ/mol 495.8 kJ/mol 122 kJ/mol -787 kJ/mol 107.3 kJ/mol -411 kJ/mol This is Hess’ Law coming in Chapt 8

46. Alkali Metals (Group IA) • Alkali metals are extremely reactive, thus they are not found free in nature. • The alkali metals have only one valence electron (ns1). • Group IA metals have low first ionization energies. • Ionization energies decrease with increasing size of the alkali metals. • All the alkali metals form stable 1+ ions. • All the IA salts are soluble (virtually no exceptions)

47. Group IIA Metals • Alkaline earth metals are silvery white, malleable, ductile, and somewhat harder than Group IA metals. • All have two valence electrons (ns2). • The ionization energies are greater for Group IIA than for Group IA metals.

48. The Halogens (Group VIIA) • The halogens get their name from the Greek halogen meaning salt formers • Like the IA and VIIIA elements, the familial similarity is pronounced within the group • All are highly reactive • Tend to react via free radical mechanism • Color of elements are due to formation of free radicals • Obtained from oxidation of halide salts Order of increasing ease of oxidation.

49. Chapter 7Covalent Bonds and Molecular Structure • Covalent Bonding • Lewis Structure • Periodic Properties • Resonanace • Polar/Nonpolar bonds • Formal Charges • Molecular Shapes • VSEPR • Dipole moment • MO Theory

50. VSEPR Theory • VSEPR dictates that the regions of high electron density around the central atom are as far apart as possible to minimize repulsions. • Key to this is that nonbonded electron pairs occupy more space than bonded electrons do. • There are five basic orbital configurations. • These shapes are based upon the number of regions of high electron density within the molecule. • Molecular shapes are based upon these basic orbital configurations.